MAIN  LIBRARY  AGRIC. 


QUALITATIVE  ANALYSIS 


QUALITATIVE  ANALYSIS 


AS    A   LABORATORY    BASIS    FOR    THE   STUDY    OF 


GENERAL  INORGANIC  CHEMISTRY 


BY 

WILLIAM  CONGER  MORGAN,  PH.D.  (YALE) 

ASSISTANT  PROFESSOR  OF  CHEMISTRY  IN  THE  UNIVERSITY  OF  CALIFORNIA 


OF   THE 

UNIVERSITY 

OF 


gotfe 

THE   MACMILLAlSr   COMPANY 
LONDON:    MACMILLAN  &  CO.,  LTD. 

1910 


COPYRIGHT,  1906 

BY  THE   MACMILLAN   COMPANY 
REPRINTED  DECEMBER,  1907;  AUGUST,  1910 


PREFACE. 


This  text  presents  the  course  of  instruction  given  by  the  author  to 
the  freshman  classes  at  the  University  of  California  where  the  exigencies 
of  vocational  courses  render  it  impracticable  for  a  large  number  of 
students  to  devote  more  than  one  year  to  general  inorganic  chemistry 
and  to  qualitative  analysis  as  well.  Therefore,  the  scheme  of  qualitative 
analysis  is  made  to  serve  as  a  means  of  correlating  the  apparently  inde- 
pendent experiments  of  general  chemistry.  Thus,  while  the  end  in 
view  is  primarily  the  imparting  of  a  knowledge  of  the  general  subject, 
familiarity  with  analytical  methods  is  obtained.  The  course  presupposes 
good  chemical  training  in  the  secondary  schools,  many  of  its  briefer 
statements  being  in  the  nature  of  a  review  rather  than  of  first  presenta- 
tion. No  attempt  has  been  made  to  emasculate  the  work  by  eliminating 
the  consideration  of  more  difficult  points.  On  the  other  hand,  the  tests 
for  individual  elements  have  been  varied  in  order  to  familiarize  the 
student  with  as  many  significant  reactions  as  possible. 

In  the  opinion  of  the  author,  instruction  in  science  should  endeavor 
to  equip  the  student  with  principles  rather  than  facts  and,  what  is  of 
still  greater  import,  it  should  train  him  in  the  use  of  these  principles 
and  in  the  application  of  them  in  explaining  the  phenomena  of  general 
experience.  The  acquisition  of  relatively  large  numbers  of  facts  or  the 
memorizing  of  printed  rules  and  laws  is  a  matter  of  comparatively  easy 
accomplishment.  The  ability  to  analyze  a  situation  and  to  recognize 
the  underlying  principle  comes  less  readily.  Still  more  difficult  of 
acquirement  is  that  mental  habit  which  consistently  seeks  the  explana- 
tion of  the  unfamiliar  by  means  of  the  well-known.  In  the  author's 
experience,  furthermore,  students  know  facts  but  lack  the  perspective 
which  gives  the  appreciation  of  relative  values.  Their  facts  are  for  the 
most  part  individual  units,  like  so  many  loose  beads  without  any  con- 
necting thread  by  means  of  which  the  whole  may  be  viewed  as  a  unit 
and  the  relation  of  the  individuals,  each  to  the  other,  recognized.  From 
certain  of  the  many  books  at  present  on  the  market  a  good  student  may 
learn  much  chemistry ;  but  the  author  has  found  it  necessary  to  supply 
in  lecture  form  the  correlating  factors  and  to  keep  pointing  out  the 
application  of  general  principles  to  cases  not  less  typical  than  the 
stereotyped  illustration  discussed  in  the  text.  With  the  hope  of  minim- 


o  /  w  \  o  o  /r\ 


VI  PEEFACE. 

izing  the  necessity  for  lecture  work,  the  present  book  appears.  By 
frequent  repetition  the  mind  acquires  its  habits  of  functioning  and 
becomes  thoroughly  familiar  with  its  data.  By  frequent  repetition, 
therefore,  the  present  text  endeavors  to  teach  the  fundamental  prin- 
ciples of  chemistry  and  to  inculcate  the  habit  of  seeking  to  explain  new 
phenomena,  as  they  occur  from  time  to  time,  by  the  methods  of  past 
experience.  It  endeavors  to  reduce  to  a  rational  basis  the  multitude  of 
chemical  reactions,  a  knowledge  of  which  ofttim'es  seems  to  the  student 
to  be  a  mere  matter  of  tedious  memorizing. 

The  metlwd  of.  presentation  is  based  principally,  though  not  ex- 
clusively, upon  the  theory  of  electrolytic  dissociation.  Equations  are 
frequently  written,  first,  in  the  older  style  expressing  whole  molecules 
and  afterwards  as  ionic  reactions.  This  plan  conforms  to  good  peda- 
gogical principles  by  passing  from  what  is  known  to  the  student  to  what 
is  unknown  to  him.  The  presentation  appears,  therefore,  as  a  statement 
of  fact  revised  and  rendered  somewhat  more  complete  than  the  older 
form  with  which  he  is  familiar. 

In  the  experience  of  the  author,  no  difficulty  in  the  presentation  of 
the  ionic  theory  has  been  encountered.  Its  essentials  seem  to  be  grasped 
as  readily  and  comprehended  as  correctly  as  is  any  other  theoretical 
conception.  No  field  offers  more  illustrations  of  ionic  reactions  than 
qualitative  analysis.  Here  the  unifying  and  systematizing  value  of  a 
fundamental  conception  seems  to  be  greatly  needed.  The  author  would 
dissent,  therefore,  from  the  opinion  of  those  who  hold  that  instruction  in 
qualitative  analysis  should  not  be  based  on  the  dissociation  theory, 
''however  valuable  this  concept  may  have  proved  to  chemistry." 

The  text  is  divided  into  four  parts.  The  general  principles  under- 
lying any  intelligent  comprehension  of  the  subject  are  presented  briefly 
in  an  introductory  part.  The  reactions  of  individual  elements  (arranged 
according  to  the  Periodic  System)  comprise  th'e  second  part.  The  pres- 
entation of  analytical  methods  constitutes  the  third  part  while  several 
tables  of  useful  information  are  found  in  an  appendix. 

It  is  hardly  necessary  to  remark  that  the  metric  system  of  weights 
and  measures  and  the  Centigrade  scale  are  used  throughout  the  text. 
In  the  analytical  section  certain  abbreviations  occur  which  will  be 
understood  without  difficulty.  The  solutions  used  in  the  laboratory  are 
of  the  concentrations  given  in  the  appendix.  Throughout  the  text, 
whenever  no  concentration  is  specified,  the  solutions  of  acids  and  alkalies 
are  understood  to  be  the  dilute  reagents,  four  times  normal. 

In  the  preparation  of  the  text  most  of  the  better  works  on  general 
inorganic  chemistry  and  on  qualitative  analysis  have  been  consulted 


PREFACE.  vii 

freely  and  statements  of  fact  taken  from  them.     On  many  points,  how- 
ever, original  sources  have  had  to  be  investigated. 

In  conclusion,  grateful  acknowledgment  is  made  of  the  advice  and 
assistance  rendered  by  several  friends  who  have  appreciably  lessened  the 
work  of  publication.  In  this  connection  Professor  F.  C.  Cottrell,  of  the 
University  of  California,  should  be  especially  mentioned  inasmuch  as  he 
has  had  the  kindness  to  read  a  portion  of  the  original  manuscript.  Of 
the  constant  assistance  of  my  wife  in  the  preparation  of  the  manuscript 
I  would  also  make  mention. 

WILLIAM  CONGER  MORGAN. 

BERKELEY,  CALIFORNIA, 
October,  1906. 


CONTENTS 


PART    I.     GENERAL. 

PAGE. 

Introduction 3 

Equations 5 

Aids  to  Chemical  Reactions 7 

Dissociation    .                  9 

Acids,  Bases  and  Salts 15 

Equilibrium  and  Mass  Action 22 

Hydrolysis 28 

Repression  of  lonization 31 

Solute  and   Solvent 32 

Separation  of  Solids  from  Liquids 36 

Oxidation  and  Reduction 38 

Preliminary  Examination .         .53 

Weights  and  Measures 57 

PART    II.    DESCRIPTIVE. 

Classification  of  the  Elements — The  Periodic  System 62 

Group  0 65 

Group  1,  First  Type 65 

Lithium 66 

Sodium    .                  67 

Potassium 72 

Ammonium .         .         .75 

Group  1,  Second  Type 79 

Copper 79 

The  Cupric  Ion 81 

The  Cuprous  Ion 83 

Silver 84 

Gold 89 

Group  2,  First  Type 90 

Magnesium 91 

Calcium 94 

Strontium 97 

Barium 98 

ix 


X  CONTENTS. 

Group  2,  Second  Type IQQ 

zinc                                    101 

Cadmium 104 

Mercury           .                  106 

The  Mercuric  Ion 107 

The  Mercurous  Ion HO 

Group  3,  First  Type     .7 m 

Boron H2 

Boric  Acid H2 

Aluminum 115 

Group  3,  Second  Type US 

Group  4,  First  Type 119 

Carbon 119 

Carbonic   Acid 121 

Acetic  Acid 122 

Oxalic  Acid .         .  123 

Tartaric  Acid 124 

Cyanogen          .                           ........  125 

Hydrocyanic  Acid 125 

Cyanic  Acid 128 

Thiocyanic  Acid 129 

Ferrocyanic  Acid .  130 

Ferricyanic  Acid 131 

Silicon 132 

Silicic  Acid .         .         .133 

Hydrosilicofluoric  Acid 136 

Group  4,  Second  Type 137 

Tin 137 

The  Stannous  Ion 139 

The  Stannic  Ion 140 

Lead 142 

Group  5,  First  Type     ... 145 

Nitrogen 146 

Nitrous  Acid 148 

Nitric   Acid 149 

Phosphorus 151 

Phosphoric   Acid 153 

Arsenic 156 

The  Arsenious  Kathion 159 

Arsenic  Compounds 160 

Arsenious  Acid 161 

Arsenic  Acid  .                           .  163 


CONTENTS.  xi 

Antimony .  165 

The  Antimonous  Ion 166 

Pentavalent  Antimony  Compounds 16S 

Antimonic  Acid 168 

Bismuth .169 

Group  5,  Second  Type 172 

Group  6,  First  Type * 172 

Oxygen 173 

Water 176 

Hydrogen  Dioxid 177 

Sulfur 179 

Hydrosulfuric  Acid 180 

Sulfurous  Acid 184 

Sulfuric  Acid 186 

Thiosulfuric  Acid 188 

Group  6,  Second  Type 190 

Chromium 190 

The  Chromic  Kathion 192 

The  Chromate  Ion  and  the  Bichromate  Ion  .         .         .         ...  194 

Group  7,  First  Type 197 

Fluorin 198 

Hydrofluoric  Acid 198 

Chlorin 200 

Hydrochloric  Acid 202 

Hypochlorous  Acid 203 

Chloric  Acid 205 

Perchloric  Acid 206 

Bromin 208 

Hydrobromic  Acid 209 

lodin 210 

Hydriodic  Acid 211 

Group  7,  Second  Type 212 

Manganese 213 

The  Manganous  Ion 216 

The  Permanganate  Ion 218 

Group  8 220 

Iron 220 

The  Ferrous  Ion 222 

The  Ferric  Ion 224 

Cobalt 227 

Nickel 231 

Platinum 234 

The  Care  of  Platinum  Ware  .  .  235 


xii  CONTENTS. 


PART   III.    ANALYTICAL. 

The  Systematic  Analysis 239 

Preliminary  Examination 243 

A.  Heat  in  Closed-tube 243 

B.  Heat  with  Concentrated  Sulfuric  Acid 245 

C.  Heat  on  Charcoal 246 

D.  Heat  in  Borax-bead 247 

E.  Heat  on  Platinum  Wire 247 

The  Basic  Analysis .  248 

Group  Separations  ...  248 

Special  Tests 249 

Separation  and  Analysis  of  Group  I,  Div.  A 250 

Discussion 251 

Separation  of  Group  I,  Divs.  B  and  C 252 

Discussion 252 

Analysis  of  Group  I,  Div.  B .  ~ .  .  254 

Discussion 255 

Analysis  of  Group  I,  Div.  C 256 

Discussion 257 

Analysis  of  Residue  10  when  Gold  or  Platinum  may  be  present  .  .  258 

Discussion 259 

Detection  of  Interfering  Substances .  260 

Discussion 260 

Removal  of  Interfering  Substances 262 

Discussion         .....         ......  263 

Separation  of  Group  II 264 

Discussion 265 

Analysis  of  Group  II  .  .  . 266 

Discussion 267 

Separation  and  Analysis  of  Group  III 268 

Discussion 269 

Separation  and  Analysis  of  Group  IV 270 

Discussion 271 

Analysis  of  Group  V 272 

Discussion 273 

The  Acid  Analysis 275 

Preliminary  Tests 276 

Discussion 277 

Test  for  Sulfuric  Acid 278 

Discussion 279 

Test  for  Hydrosilicofluoric  Acid 278 

Discussion 279 

Tests  for  Hydrofluoric  Acid 278 

Discussion        .        .  279 


CONTENTS.  xni 

Tests  for  Silicic  Acid .280 

Discussion 281 

Test  for  Dichromic  Acid 280 

Discussion 281 

Tests  for  Boric  Acid 282 

Discussion 283 

Test  for  Carbonic  Acid "282 

Discussion 283 

Test  for  Oxalic  Acid .284 

Discussion 285 

Tests  for  Tartaric  Acid  .  .  .  284 

Discussion 285 

Tests  for  Phosphoric  Acid  ...  .  .  286 

Discussion 287 

Test  for  Arsenic  Acid 288 

Discussion 289 

Test  for  Arsenious  Acid •  288 

Discussion 289 

Test  for  Sulfurous  Acid  .  ...  .290 

Discussion 291 

Test  for  Thiosulfuric  Acid  ....  ....  290 

Discussion 291 

Test  for  Ferrocyanic  Acid 292 

Discussion  .  .  .  .  . 293 

Test  for  Ferricyanic  Acid 292 

Discussion 293 

Tests  for  Thiocyanic  Acid 292 

Discussion 293 

Tests  for  Hydrocyanic  Acid 294 

Discussion 295 

Tests  for  Hydrosulfuric  Acid 294 

Discussion 295 

Tests  for  Hydriodic  Acid 296 

Discussion         ...........  297 

Tests  for  Hydrobromic  Acid 298 

Discussion 299 

Tests  for  Hydrochloric  Acid 298 

Discussion 299 

Tests  for  Nitric  Acid 300 

Discussion 301 

Tests  for  Nitrous  Acid .  302 

Discussion 303 

Tests  for  Acetic  Acid 302 

Discussion 303 


xiv  CONTENTS. 

Tests  for  Cyanic  Acid- 

Discussion 
Test  for  Hypochlorous  Acid 

Discussion 
Tests  for  Chloric  Acid 

Discussion         .         .         .         . 
Test  for  Perchloric  Acid 

Discussion 
Test  for  Permanganic  Acid 

Discussion         .  .         .  .. 

Tests  for  Hydrogen  Peroxid 

Discussion 
Anatysis  of  Organic   Substances 

Discussion 

APPENDIX. 

Reagents  in  Solution     .         .         .         .     •    . 

Acids 

Alkalies 

Other  Reagents  in  Solution  or  in  Liquid  Form 

Salt   Solutions 

Reagents  Requiring  Special  Methods  of  Preparation 
Reagents  in   Solid  Form 
Specific  Gravity  of  Solutions  at  15°   .         .         .         . 

Sulfuric  Acid 

Hydrochloric  Acid  .         .         .         .         . 

Nitric    Acid 

Ammonia  Solutions 
List  of  Elements 
Nomenclature 

Physical  Properties  of  Elements 
Potential   Series     ....  ... 

Metric  Measures  with  English  Equivalents  ... 
English  Measures  with  Metric  Equivalents  ... 
Conversion  Factors 
Table  of  Solubilities 


QUALITATIVE  ANALYSIS 


PART  I 


GENERAL 


QUALITATIVE  ANALYSIS 


Introduction. 

Chemistry  is  that  part  of  natural  science  treating  of  the  composition 
of  substances.  Knowledge  of  the  composition  of  substances  is  obtained 
by  three  general  methods,  analysis,  synthesis,  and  metathesis. 

Analysis  is  the  taking  apart  or  breaking  down  of  a  substance  into 
simpler  substances,  as : 

FeS  ->  Fe  +  S        CaC03  ->  CaO  +  C02 

Synthesis  is  the  putting  together  and  building  up  "of  more  complex 
compounds,  as: 

Fe  +  S  -»  FeS        CaO  +  C02  ->  CaC03 

Metathesis  is  a  process  involving  the  simultaneous  decomposition 
of  two  substances  and  the  formation  of  two  new  bodies  from  their  com- 
ponents, as: 


NaCl  +  AgN03  -»  AgCl  +  NaNO 


The  general  term  chemical  analysis  includes  all  those  processes  of 
whatever  nature  by  virtue  of  which  we  are  enabled  to  ascertain  the 
composition  of  a  body.  They  are  generally  of  an  analytical  or  meta- 
thetical  character,  but  sometimes  valuable  information  is  furnished  by 
synthetical  methods. 

Chemical  analysis  is  divided  into  two  parts,  qualitative  and  quanti- 
tative. Qualitative  analysis  is  concerned  with  the  kinds  of  matter 
composing  a  given  substance  without  regard  to  the  quantities  of  the 
constituents  present. 

Quantitative  analysis  determines  the  proportions  in  which  the  con- 
stituents are  present.  A  thorough  knowledge  of  qualitative  analysis 

*  Throughout  the  text,  hereafter,  the  physical  condition  of  reagents  under  the 
conditions  of  the  experiment  will  be  indicated  whenever  equations  are  written.  Solids 
or  insoluble  substances  will  be  denoted  by  black-faced  type,  liquids  or  substances  in 
solution  by  ordinary  type  and  gases  by  italics. 

3 


4  QUALITATIVE    ANALYSIS. 

is  prerequisite  to  the  satisfactory  understanding  of  quantitative  methods. 

Aside  from  information  as  to  the  composition  of  the  various  bodies 
about  us,  qualitative  analysis  affords  a  very  convenient  method  of  co- 
ordinating many  facts  of  general  chemistry,  for  by  a  well-arranged 
series  of  the  apparently  independent  experiments  of  general  chemistry 
we  are  enabled  to  develop  analytical  methods.  By  the  frequent  labora- 
tory repetition  and  thoughtful  consideration  of  these  processes,  the  facts 
of  general  chemistry  become  fixed  in  the  mind. 

Pure  substances  possess  characteristic  properties  which  are  un- 
changing. Chief  among  these,  for  the  present  purpose,  are:  physical 
condition  at  ordinary  temperatures,  viz.,  solid,  liquid,  or  gas;  degree  of 
solubility;  color;  physical  structure;  specific  gravity;  melting-point 
and  others  of  similar  nature.  By  a  combination  of  a  sufficient  number 
of  properties  any  substance  may  be  differentiated  from  all  others,  and  it 
is  by  virtue  of  its  properties  that  we  know  and  recognize  each  substance. 
It  is  important,  therefore,  to  observe  carefully  and  to  make  a  study  of 
the  physical  properties  of  the  substances  with  which  one  comes  in 
contact.  It  is  only  by  noting  the  properties  of  the  various  products 
which  may  be  obtained  during  a  qualitative  analysis,  by  comparing  these 
properties  with  the  properties  of  known  substances  and  realizing  their 
identity  or  difference  that  one  is  able  to  determine  the  nature  of  a 
body  analyzed. 

Mixtures  do  not  ordinarily  possess  characteristic  properties  by  which 
they  may  be  recognized.  While  the  properties  of  mixtures  are  usually 
a  blending  of  the  properties  of  the  constituents,  this  is  not  always  so. 
Frequently  mixtures  do  not  possess  the  properties  common  to  the  constitu- 
ents in  an  average  degree,  but  certain  properties  are  emphasized,  others 
much  reduced  by  the  process  of  mixing.  For  this  reason  it  is  sometimes 
necessary  to  separate  the  constituents  of  a  mixture  from  each  other  and 
obtain  each  in  a  state  of  purity  before  it  can  be  recognized.  This 
is,  however,  not  usually  necessary  in  qualitative  analysis. 

Analytical  chemistry  has  so  far  succeeded  in  reducing  the  many 
varied  substances  which  constitute  the  world  around  us  into  about 
seventy-five  simpler  forms  of  matter,  called  elements,  which  may  be 
readily  differentiated,  one  from  the  other.  These  elements  show  a  re- 
markable persistence.  Most  of  them  unite  with  other  elements  to  form 
large  numbers  of  compounds,  each  one  of  which  possesses  entirely  dis- 
tinct properties.  By  appropriate  analytical  methods,  however,  the 
simple  forms  may  be  obtained  again  and  their  identity  recognized. 

Some  elements  are  found  widely  distributed  and  in  considerable 


EQUATIONS.  5 

quantities.  Others  are  exceedingly  rare.  To  develop  and  become  fa- 
miliar with  a  system  of  qualitative  analysis  that  should  include  all  the 
elements  would  be  a  task  of  extreme  difficulty.  It  would  be  of  com- 
paratively little  value,  moreover,  because  of  the  great  rarity  of  nearly 
one-half  of  the  elementary  substances.  Hence  it  is  customary  to  con- 
sider only  the  commoner  elements.  Those  included  in  the  present 
scheme  of  analysis  have  been  indicated  by  black-faced  type  in  the  list 
given  on  page  324.  The  letter  or  letters  following  each  name  is  the 
symbol  by  which  the  element  is  generally  represented. 

Qualitative  analysis  consists  in  determining  primarily  what  elements 
are  contained  in  any  given  substance.  Many  of  the  elements  are  so 
reactive,  however,  that  they  are  very  rarely  found  uncombined  with 
others.  Furthermore,  they  are  not  separated  in  the  free  state  during 
the  processes  of  analysis.  On  the  other  hand,  many  combinations  of  two 
or  more  elements,  fairly  stable  under  definite  conditions,  are  of  frequent 
occurrence  and,  since  these  are  well  recognized,  it  is  generally  not 
necessary  to  separate  a  body  into  its  elementary  constituents,  but  is 
often  sufficient  to  determine  the  presence  of  the  groups,  or  radicles. 

The  method  of  chemical  analysis  is  to  change  an  unknown  substance 
into  other  substances,  the  properties  and  composition  of  which  are  well 
known.  From  a  consideration  of  the  derived  substances  judgments  are 
formed  as  to  the  composition  of  the  original  substance.  Such  a  change 
is  known  as  a  reaction.  It  is  frequently  accomplished  by  the  addition  of 
a  substance  of  known  composition  and  properties  called  a  reagent. 

Reactions  carried  on  without  the  presence  of  appreciable  quantities 
of  water  are  said  to  be  performed  in  the  ' '  dry  way, ' '  those  taking  place 
in  aqueous  solution  in  the  "  wet  way." 

EQUATIONS. 

An  equation  is  a  chemical  short-hand  expression  of  the  changes 
taking  place  during  a  reaction,  the  symbols  for  the  elements  and  their 
compounds  being  used.  All  the  substances  which  disappear  wholly  or 
in  part  during  the  reaction  are  written  on  one  side,  while  all  the  sub- 
stances formed  during  the  reaction  are  grouped  on  the  other  side,  sep- 
arated from  the  first  group  by  an  arrow  indicating  the  direction  in  which 
the  reaction  proceeds,  thus: 

Na  +  I  ->  Nal 

For  general  use  the  arrow  is  to  be  preferred  to  the  sign  of  equality, 
for  the  equation  is  not  ordinarily  considered  as  indicating  that  one  side 


6  QUALITATIVE    ANALYSIS. 

is  equal  to  the  other  side,  but  only  that  one  group  of  reagents  passes 
over  into  the  other. 

In  writing  formulae  for  chemical  compounds  or  equations  for  reac- 
tions the  symbol  of  an  element  stands  for  a  definite  quantity  of  that 
element,  for  one  atom*  only.  Thus  H  stands  for  one  atom  of  hydrogen 
and  0  for  one  atom  of  oxygen.  If  more  than  one  atom  is  present,  the 
fact  is  indicated  by  a  small  number  written  subscript,  as  H2.  Such  a 
subscript  number  affects  the  symbol  only  which  it  immediately  follows. 
Thus,  H20  means  two  atoms  of  hydrogen  and  one  of  oxygen.  The 
formula  of  a  compound  expresses  the  composition  of  a  molecule*  of  the 
substance.  The  symbols  of  the  elements  composing  it  are  written  to- 
gether, the  proper  number  of  atoms  being  indicated  by  subscript  num- 
bers, as  NaCl,  H20,  and  K2O207.  If  it  be  desired  to  indicate  more  than 
one  molecule  of  a  compound,  a  large  number  is  prefixed  to  the  formula. 
This  acts  as  an  algebraic  coefficient  and  multiplies  each  element  in  the 
formula.  Two  molecules  of  sodium  chloride  or  potassium  carbonate 
are  written:  2NaCl  and  2K2CO3,  and  not  Na2Cl2  or  K4C206. 

All  equations  should  be  balanced.  The  total  amount  of  matter  which 
appears  on  one  side  should  be  present  in  some  form  on  the  other,  the 
same  number  of  atoms  of  each  element  appearing  on  each  side,  as  is 
shown  in  the  following  equation : 

2KMn04  +  3H2S04  +  10HC1  ->  K2S04  +  2MnS04  +  8H20  +  IOCI 

In  order  to  do  this  it  is  necessary  to  resort  to  quantitative  methods 
and  determine  the  quantities  of  the  reagents  which  enter  into  a  reac- 
tion. These  determinations  have  been  satisfactorily  made  and  need  not 
be  undertaken  by  the  student,  but  a  knowledge  of  the  quantitative  rela- 
tions in  any  given  equation  may  be  ascertained  by  reference  to  proper 
sources  or  by  comparison  with  the  known  values  in  similar  reactions. 

A  chemical  equation  should  be  a  statement  of  fact.  Before  a  state- 
ment of  fact  can  be  made  it  is  necessary  to  know  the  facts.  Hence  it  is 
necessary  to  know  what  actually  takes  place  in  a  given  reaction  before 
the  equation  can  be  correctly  written.  Chemical  reactions  take  place  for 
the  most  part  according  to  well-established  laws.  A  familiarity  with 
these  laws  and  with  the  nature  of  chemical  elements  and  compounds 
will  often  enable  one  to  tell  at  a  glance  from  the  nature  of  the  sub- 
stances present  and  the  conditions  of  the  experiment  what  reaction  will 
take  place.  It  is  essential  to  a  good  understanding  of  the  subject  that 

*  For  the  definitions  of  the  terms  atom  and  molecule  see  page  8. 


AIDS    TO    CHEMICAL    EEACTIONS.  7 

the  student  be  able  to  write  equations  correctly  expressing  the  changes 
taking  place  in  the  experiments  which  he  is  performing.  In  general  the 
ability  to  write  chemical  equations  may  be  taken  as  an  index  of  the 
chemical  knowledge  possessed  by  the  student  and  of  his  understanding 
of  the  processes  which  he  is  performing. 

AIDS    TO    CHEMICAL    REACTIONS. 

Solids  do  not  readily  react  with  each  other  under  ordinary  con- 
ditions. Liquids  and  gases  are  more  reactive,  however.  In  order  to 
cause  a  substance  under  examination  to  react,  it  is  frequently  necessary 
to  aid  the  reaction  by  the  application  of  (a)  mechanical  force,  (6)  heat, 
(c)  light,  (d)  electricity  and  (e)  solution. 

(a)  As  an  illustration  of  the  action  of  mechanical  force  it  may  be 
stated  that  potassium  chlorate  and  sulfur,  however  finely  powdered 
and  intimately  mixed,  will  not  react.  They  may  be  made  to  unite  with 
explosive  violence*  by  pressure  or  friction. 

(6)  By  the  application  of  heat  solids  ordinarily  melt  and  become 
liquids.  If  the  temperature  be  raised  still  higher,  liquids  volatilize  and 
become  gases. 

Aside  from  the  formation  of  liquids  and  gases  which  are  more 
reactive  than  solids,  heat  aids  chemical  activity  without  causing  a  change 
in  the  physical  state.  Many  substances,  which,  at  ordinary  tempera- 
tures, will  not  react  to  any  appreciable  degree,  will  do  so  with  the 
greatest  readiness  at  higher  temperatures.  A  fairly  definite  relation 
between  the  increase  in  temperature  and  the  time  necessary  to  complete 
a  reaction  has  been  found  to  exist.  As  the  average  result  of  experimental 
investigation  the  general  statement  may  be  made  that  the  speed  of  a 
reaction  is  doubled  by  every  increase  of  10°  Centigrade.  Thus  by 
raising  the  temperature  100°  the  reaction  would  be  hastened  more 
than  a  thousand  fold,  and  a  reaction  which  would  not  take  place  ap- 
preciably during  an  interval  of  years  at  ordinary  temperatures  might 
be  accomplished  in  a  few  minutes  at  1000°.  Hence  it  is  assumed  that 
all  reactions  that  take  place  at  high  temperatures  do  actually  occur 
at  ordinary  temperatures,  but  with  a  reaction  speed  which  approximates 
zero.  Heat  as  an  agent  to  hasten  chemical  activity  is  of  very  wide 
application,  and  frequently  productive  of  very  marked  results. 

(c)  To  a  much  smaller  degree  light  aids  chemical  reactions,  photo- 
chemical action,  as  it  is  called,  being  the  foundation  of  all  photographic 
processes. 

*  Undoubtedly  most  reactions  of  this  class  may  be  explained  as  due  to  the  action 
of  heat  generated  by  mechanical  force. 


8  QUALITATIVE    ANALYSIS. 

(d)  Electricity  is  sometimes  called  in  to  aid  a  chemical  reaction  as 
in  the  production  of  pure  copper.     The  crude  "matte"  is  dissolved  and 
the  resulting  copper  sulphate  solution  electrolyzed  with  the  separation 
of  the  metal. 

(e)  A  large  proportion  of  all  substances  known  to  interact  under 
any  conditions  will  do  so  readily  in  solution,  especially  in  aqueous  solu- 
tion.   Many  of  these  reactions  cannot  be  made  to  take  place  under  any 
other  condition.     Solution  in  water  is,  therefore,  one  of  the  greatest 
known  aids  to  chemical  activity. 

According  to  present  theories,  matter  is  composed  of  minute  par- 
ticles in  constant  motion.  These  molecules,  as  they  are  called,  are 
denned  as  the  smallest  particles  of  any  substance  that  can  exist  and  pos- 
sess the  properties  of  that  substance.  Further,  each  molecule  of  a 
compound  substance  may  be  divided  into  the  still  smaller  particles  of 
the  elements  composing  it.  When  so  divided,  however,  the  properties 
of  the  compound  substance  are  no  longer  present,  but  only  the  proper- 
ties of  a  mixture  of  elements.  These  minute  particles  of  the  elementary 
substances  which  act  as  units  in  the  formation  of  chemical  compounds 
are  called  atoms.  Atoms  are  supposed  to  be  in  constant  motion  within 
the  molecule. 

Heat  and  light,  so  far  as  their  nature  is  known,  and  possibly  elec- 
tricity, seem  to  be  of  a  vibratory  character.  Their  action  in  causing 
increased  chemical  activity  may  easily  be  explained  as  the  result  of  a 
greater  freedom  of  motion  given  to  the  atoms  and  molecules.  There  is, 
hence,  a  greater  tendency  for  the  molecules  to  break  down  and  for 
the  atoms  to  enter  into  new  combinations  which  are  more  stable  under 
the  conditions  present. 

The  greater  freedom  of  motion  of  the  different  parts  within  a 
chemical  compound  thus  indicated  theoretically  may  be  demonstrated 
practically.  If  ammonium  chlorid  be  heated  strongly,  it  breaks  down, 
partially  at  least,  into  hydrogen  chlorid  and  ammonia,  which  may  be 
separated  and  identified  by  appropriate  means.  On  lowering  the  tem- 
perature the  two  parts  come  together  again  forming  the  original  sub- 
stance. Such  a  decomposition  is  known  as  dissociation.  A  reaction 
which  may  go  in  either  direction  is  said  to  be  reversible.  This  is 
indicated  by  a  double-headed  arrow,  as  in  the  following  equation: 


DISSOCIATION.  9 

DISSOCIATION. 

If  heat,  light  and  electricity  aid  chemical  activity  by  increasing  the 
motion  within  the  molecules,  it  is  not  improbable  that  the  increased 
activity  of  substances  in  aqueous  solution  may  be  explained  similarly 
on  the  hypothesis  of  a  greater  freedom  imparted  to  the  components  of 
the  molecules.  There  is  abundant  evidence  to  indicate  that  in  solution 
the  dissolved  substance  exists  in  a  condition  so  much  like  the  gaseous 
state  that  the  same  laws  can  be  made  to  apply  to  a  large  extent.  If 
dissociation  be  a  phenomenon  of  the  gaseous  state  may  not  some  similar 
separation  of  the  components  of  the  molecules  take  place  in  solution? 
There  is  evidence  that  such  is  the  case.  Substances  whose  activity  is 
increased  by  dissolving  them  in  water  dissociate  on  going  into  solution. 

The  results  of  the  process  of  dissociation  in  aqueous  solution  are 
very  analogous  to  the  results  of  dissociation  in  the  gaseous  state.  The 
process,  however,  seems  to  be  more  complex.  Evidence  is  accumulating 
to  indicate  that  the  solvent  unites  with  the  components  of  the  molecules 
after  they  have  dissociated  in  a  way  which  is  not,  as  yet,  clearly  under- 
stood. In  the  present  considerations  it  will  not  be  necessary  to  em- 
phasize this  phase  of  the  process  since  the  components  of  the  molecules 
act  in  many  respects  as  though  they  were  not  associated  with  the 
solvent  to  a  degree  sufficient  to  influence  chemical  reactions. 

When  common  salt  dissolves  in  water,  it  dissociates  more  or  less  com- 
pletely into  minute  particles  of  sodium  and  chlorin  called  ions  and,  for 
the  most  part,  does  not  exist  as  molecules  in  which  the  metal  is  united 
with  the  non-metallic  element.  The  process  is  known  as  ionization  or 
dissociation. 

NaCl^Na  +  Cl 

That  these  ions,  always  present  in  any  ordinary  salt  solution,  possess 
very  different  properties  from  metallic  sodium  and  gaseous  chlorin 
must  be  evident  to  anyone  familiar  with  these  substances.  Metallic 
sodium  ordinarily  decomposes  water,  liberating  hydrogen  and  forming 
sodium  hydroxid.  Gaseous  chlorin  normally  reacts  with  water,  liber- 
ating oxygen,  and  forms  hydrochloric  acid.  It  also  readily  attacks  so- 
dium, forming  sodium  chlorid,  yet  the  ions  of  sodium  and  chlorin  exist 
side  by  side  in  solution  without  appreciable  decomposition  of  the  solvent 
in  the  usual  way. 

The  difference  seems  to  be  due  mainly  to  the  fact  that  each  ion  is 
charged  with  electricity.  That  this  is  so  and  that  the  electrical  charge 
accounts  for  the  difference  in  properties  is  indicated  by  the  following 


10  QUALITATIVE    ANALYSIS. 

facts:  Two  pieces  of  platinum  may  be  charged  with  electricity,  the  one 
positively  and  the  other  negatively,  by  being  connected  with  a  cell,  dynamo 
or  other  source  of  electricity.  When  so  connected  they  are  known  as 
electrodes.  If  these  electrodes  be  immersed  in  a  solution  made  by  dis- 
solving sodium  chlorid  in  pure  water,  each  loses  its  charge,  although 
it  is  a  well-known  fact  that  pure  water  is  not  an  appreciable  conductor  of 
electricity.  Of  course,  the  charges  on  the  platinum  strips  are  renewed 
by  the  cell  or  dynamo  and,  by  suitably  modifying  the  apparatus,  sodium 
may  be  collected  on  one  electrode  while  chlorin  is  liberated  on  the  other. 

The  explanation  seems  to  be  as  follows:  The  sodium  ions  are  at- 
tracted to  the  electrode  charged  with  negative  electricity  because  they 
carry  a  positive  charge  of  electricity.  At  any  rate,  they  leave  the  solu- 
tion entirely  and  deposit  on  the  electrode,  which  loses  its  charge  of 
negative  electricity  exactly  as  if  neutralized  by  the  addition  of  plus 
electricity.  But  as  soon  as  the  sodium  ions  give  up  their  charges,  they 
take  on  all  the  properties  of  metallic  sodium  and,  unless  protected  from 
the  action  of  the  water,  they  react  with  it  in  the  usual  way. 

Similarly  the  chlorin  ions,  being  negatively  charged,  deposit  on  the 
positively  charged  electrode.  They  give  up  their  electricity,  cease  to 
be  ions,  and  assume  all  the  properties  of  gaseous  chlorin.  In  this  way 
the  process  of  electrolysis,  or  analysis  ~by  means  of  an  electric  current, 
is  explained.  No  solution  will  conduct  an  electric  current  that  does  not 
contain  ions.  This  decomposition,  by  virtue  of  which  electrolysis  becomes 
possible,  is  sometimes  called  electrolytic  dissociation,  and  the  substance 
that  dissociates,  the  electrolyte. 

The  ions  formed  ~by  the  metals  and  ~by  hydrogen  are  charged  with 
plus  electricity  and  in  electrolysis  go  to  the  kathode  or  negative  electrode. 
They  are  called  kathions.  Non-metallic  ions,  charged  with  minus  elec- 
tricity, go  to  the  anode  or  positive  electrode  and  are  called  anions. 

The  degree  of  dissociation  increases  as  the  dilution  increases.  The 
greater  the  volume  of  water  in  which  a  given  amount  of  a  substance  is 
dissolved,  the  more  completely  will  it  be  ionized.  Conversely,  by  decreas- 
ing the  volume  of  the  solvent,  dissociation  is  lessened,  for  the  ions  come 
together  to  form  undissociated  substance. 

It  may  be  demonstrated  readily  that  no  electricity  passes  off  with 
the  vapors  of  the  solvent  upon  evaporation.  If  a  salt  solution,  com- 
pletely dissociated  into  sodium  and  chlorin  ions  charged  with  electricity 
of  opposite  sign,  be  evaporated  without  loss  of  electricity,  it  leaves  an 
electrically  neutral  solid.  The  quantity  of  electricity  on  the  sodium 
ions  must  be  equal,  therefore,  to  the  quantity  of  electricity  on  the 


DISSOCIATION.  11 

chlorin  ions.  Since  the  number  of  the  two  kinds  of  ions  is  equal,  the 
charge  on  each  sodium  ion  must  be  equal  in  quantity,  though  of  opposite 
sign,  to  the  charge  on  each  chlorin  ion.  In  this  way  it  may  be  demon- 
strated that  the  amount  of  electricity  on  kathions  of  any  kind  is  equal 
to  the  charge  on  the  anions  produced  with  them. 

In  a  solution  of  calcium  chlorid,  however,  there  are  two  chlorin  ions 
for  one  calcium  ion.  It  follows,  then,  that  each  calcium  ion  carries 
twice  the  charge  on  each  chlorin  ion,  or  twice  the  charge  on  each  sodium 
ion.  The  amount  of  electricity  carried  l>y  each  sodium  or  chlorin  ion 
is  taken  as  the  unit  charge.  The  calcium  ion,  therefore,  carries  two 
unit  charges. 

The  separation  of  equal  amounts  of  opposite  electricity  is  an  essential 
factor  in  the  dissociation  of  any  substance,  and  is  usually  indicated 
by  the  use  of  plus  (-(-)  and  minus  (  —  )  signs  written  above  the  symbol 
for  the  kathion  or  anion,  one  sign  being  affixed  for  each  unit  charge  of 
electricity  on  the  ion.  The  reaction  for  the  ionization  of  sodium  chlorid 
may  be  more  correctly  written  as  follows: 

NaCl  <r*  Na  +  Cl 
Similarly  for  calcium  chlorid  : 

CaCl2  <->  Ca  +  2C1 
By  anology  one  would  say: 


and 

SnCl4  <r*  Sn  +  4C1 

Ions  carrying  one  unit  charge  are  called  univalent;  ions  with  two 
charges,  divalent;  three  charges,  trivalent;  and  four  charges,  tetra- 
valent.  While  univalent  and  divalent  kathions  are  of  frequent  occur- 
rence, trivalent  kathions  are  less  common,  and  tetravalent  kathions  are 
but  little  known.  Elementary  anions  with  a  valence  of  more  than  two 
are  unknown.  (See  page  13.)  As  the  valence  rises,  the  tendency  toward 
ion  formation  decreases  while  other  influences  tending  to  remove  ions 
from  solution  increase.  (See  page  46.) 

On  evaporation  of  a  solution  containing  various  ions,  two  univalent 
anions  unite  with  one  divalent  kathion  to  form  a  solid.  Similarly,  three 
univalent  ions  unite  with  one  trivalent  ion  oppositely  charged,  and  four 
univalent  ions  with  one  tetravalent  ion.  Two  divalent  ions  unite  with 


12  QUALITATIVE    ANALYSIS. 

one  tetravalent  ion  and  three  divalent  ions  with  two  trivalent  ions. 
Consequently  the  idea  of  valence  is  extended  from  ions  in  solution  to 
substances  out  of  solution  whether  they  may  ionize  or  not.  A  univalent 
element  is  one  which  will  form  univalent  ions  or  will  unite  with  elements 
forming  univalent  ions,  atom  for  atom.  A  divalent  element  is  one  which 
will  form  divalent  ions  or  will  unite  with  elements  forming  divalent 
ions,  atom  for  atom.  Trivalent  and  tetravalent  elements  are  similarly 
define. 

In  writing  the  formulae  for  compounds  which  are  not  represented 
in  the  ionic  condition,  the  valence  of  each  element  may  be  indicated  by 
bonds  joining  the  different  elements  together,  one  bond  for  each  va- 
lence shown. 

Cl  Cl     Fe=0  Cl  0 

/  /  \  /  / 

Na  — Cl     Ca  Ca  =  0     Fe— Cl  O     Sn— Cl     Sn 

\  \  /        \\  \ 

Cl  Cl     Fe=0         \C1  0 

Cl 

Sodium  and  chlorin  are  here  indicated  as  univalent  elements,  calcium 
and  oxygen  as  divalent  elements,  iron  trivalent,  and  tin  tetravalent. 
The  valences  ordinarily  shown  by  the  elements  as  they  are  encountered 
in  qualitative  analysis  are  given  on  page  324.  Under  varying  conditions 
the  valence  of  an  element  may  change.  This  point  will  be  considered 
later.  (See  page  45.) 

Ions  need  not  necessarily  be  elementary  in  character,  but  are  often 
of  a  compound  nature  consisting  of  a  group  or  radicle  acting  as  a  unit, 
as  the  following  dissociation  equations  indicate : 

Na2C03  ->  2Na  +  C03  Ca  (N08)  2  -»  Ca  +  23*b8 

(NHJ  2S04  ->  2NH4  +  SO4 

Ions  of  a  still  more   complex   nature  may  be   formed   by   definite 

chemical   reactions   between   two    different   simple   ions   or   between    a 

+ 

simple  ion  and  an  undissociated  substance.  The  silver  ion,  Ag,  unites 
directly  with  ammonia  to  form  a  complex  ion,  Ag(NH3)2,  which  possesses 
properties  entirely  different  from  the  simple  silver  ion.  The  cyanid 
ion,  CN,  unites  in  an  analogous  way  with  silver  cyanid,  AgCN,  to  form 
a  new  ion,  Ag(CN)2,  in  which  the  metal,  silver,  constitutes  a  part  of  the 
negatively  charged  anion  instead  of  acting  as  a  simple  positively  charged 
kathion.  Ammonia  and  cyanogen  compounds  show  especial  tendencies 
to  the  formation  of  complex  ions. 


DISSOCIATION.  13 

Complex  ions  vary  very  greatly  in  the  readiness  with  which  they 
are  formed  and  in  stability  after  being  formed.  In  order  to  produce 
certain  ones  it  is  only  necessary  to  bring  together  the  reagents  necessary 
to  form  them  when  an  immediate  reaction  takes  place.  In  other  cases 
it  is  necessary  to  aid  the  reactions  in  various  ways.  There  is  always  a 
tendency  for  complex  ions  to  decompose  into  simpler  ones.  Because  of 
this  certain  solutions  contain  a  mixture  of  simple  and  complex  ions.  In 
other  cases  the  stability  of  the  complex  ion  is  so  great  that  no  appreciable 
amount  of  the  simple  ions  can  be  detected. 

Complex  anions  with  a  valence  of  three  and  four  are  met  with  in 
qualitative  analysis.  (See  page  11.) 

Water  is  the  only  one  of  the  common  solvents  causing  dissociation 
to  any  great  extent,  although  acetic  acid  and  alcohol  have  slight  dis- 
sociating properties.  Many  less  common  solvents,  such  as  liquid  hydro- 
gen cyanid,  sulfur  dioxid  and  ammonia,  and  many  fused  salts,  cause  dis- 
sociation to  a  high  degree. 

All  substances  are  not  dissociated  in  water  solutions.  Many  of  the 
compounds  of  carbon  with  hydrogen  and  other  elements,  known  as 
organic  compounds,  do  not  dissociate  at  all.  On  the  other  hand,  most  of 
the  soluble  mineral  or  inorganic  substances  are  dissociated.  These  are 
the  substances  with  which  qualitative  analysis  is  chiefly  concerned. 

This  fact  greatly  simplifies  qualitative  analysis.  Each  soluble  salt 
of  a  given  metal  with  the  various  acids  gives  the  same  kathion.  Simi- 
larly, each  soluble  salt  of  a  given  acid  with  the  various  metals  gives 
the  same  anion.  It  makes  no  difference  whether  calcium  nitrate, 
chlorid,  acetate,  or  bromid  is  used ;  any  soluble  calcium  salt  will  give  the 
same  kathion.  It  is  a  characteristic  of  the  calcium  ion  that  it  gives  a 
white  precipitate  with  a  solution  of  sodium  carbonate,  or  with  the 
solution  of  any  other  soluble  carbonate  containing  the  carbonate  ion  in 
appreciable  quantities.  Hence  it  is  not  necessary  to  learn  the  reaction 
of  calcium  chlorid  on  sodium  carbonate,  of  calcium  nitrate  on  ammonium 
carbonate  and  of  calcium  sulfate  on  potassium  carbonate,  but  simply 
to  study  the  various  ions  and  learn  the  characteristic  reactions  of  each. 

The  reactions  used  in  qualitative  analysis  are  almost  entirely 
reactions  between  ions.  This  is  well  illustrated  by  the  following  facts: 
It  is  loosely  stated  that  silver  nitrate  is  a  reagent  for  chlorin,  forming 
with  it  an  insoluble  white  compound.  Silver  nitrate  is  a  reagent  for 
the  chlorin  ion  only.  Most  organic  compounds  containing  chlorin,  even 
chloroform  which  is  almost  ninety  per  cent,  chlorin,  give  no  precipitate 
with  silver  nitrate  because  they  do  not  dissociate.  Chlorates  and  per- 


14  QUALITATIVE    ANALYSIS. 

chlorates,  although  they  ionize  readily,  give  no  precipitate  because  the 
ions  formed  are  C103  and  C1O4  and  not  the  simple  Cl.  Free  chlorin 
gives  not  a  trace  of  a  precipitate  as  long  as  it  remains  un-ionized.  On  the 
other  hand,  every  substance  that  dissociates  into  Cl  ions  gives  an  abun- 
dant precipitate.  Thus  there  is  ample  evidence  that  it  is  the  chlorin 
ion  alone  that  reacts  with  silver  nitrate.  More  correctly  speaking, 
it  is  the  silver  ion,  Ag,  and  not  silver  nitrate  that  reacts  with  the  Cl 

ion,   for  undissociated  silver  nitrate   does  not  react,   neither  will  the 

+ 

ions  Ag(NH3)2  or  Ag(CN)2  or  any  other  complex  ion  containing  silver. 
Any  compound  forming  the  simple  silver  ion  will  react,  however.  Since 
silver  nitrate  is  the  only  common  silver  compound  readily  soluble,  it  is 
almost  universally  used  to  introduce  the  silver  ion,  and,  hence,  silver 
nitrate  is  often  mentioned  when  the  silver  ion  is  really  meant. 

Dissociation  is  the  characteristic  phenomenon  of  dilute  solutions. 
In  concentrated  solutions,  not  only  is  dissociation  reduced,  but  substances 
tend  to  associate  or  come  together  in  various  ways.  The  principles  and 
laws  which  hold  for  dilute  solutions  are  only  partially  true  for  con- 
centrated solutions  and  a  different  set  of  phenomena  are  to  be  observed. 
One  of  the  most  important  types  of  association  is  the  formation  of 
double  salts,  such  as  the  double  sulfates  of  aluminum  (and  certain 
other  trivalent  metals)  with  the  alkali  metals.  As  a  typical  illustration 
ordinary  alum,  KA1(S04)2,  or  as  the  formula  is  often  written, 
K2S04.A12(S04)3,  may  be  taken.  In  a  dilute  solution  containing  the 
sulfates  of  potassium  and  of  aluminum,  only  simple  ions  such  as  K, 
Al  and  S04  are  present.  As  the  solution  concentrates,  these  ions  come 
together  to  form  the  undissociated  sulfates  and,  on  further  evaporation 
of  the  solvent,  these  simple  salts  unite  to  form  the  double  salt  which 
crystallizes  out. 

Double  salts  are  to  be  regarded  as  salts  of  complex  acids  dissociating 
primarily  into  complex  anions  which,  being  very  unstable,  break  down 
into  simple  ions  as  soon  as  the  dilution  becomes  sufficiently  great  to  allow 
ionization  to  become  the  dominant  phenomenon.  The  more  stable  of  the 
double  salts  are  those  whose  anions  do  not  decompose  into  simple  ions 
readily  but  which  even  in  dilute  solutions  give  characteristic  complex 
ions.  The  double  cyanids  of  iron  with  other  metals  are  of  this  kind. 
(See  page  223.) 

The  following  equation  wrill  serve  to  make  clear  the  way  in  which  a 
double  salt  dissociates : 


ACIDS.  15 

The  complex  ion,  A1(S04)2,  is  not  known  with  certainty  and  pos- 
sibly has  only  a  momentary  existence  but  all  degrees  of  stability  from  this 
up  to  the  most  stable  are  well  recognized. 

ACIDS,  BASES  AND  SALTS. 

A  study  of  the  various  chemical  compounds  has  revealed  the  fact 
that  many  substances,  apparently  differing  as  widely  in  chemical  nature 
as  they  do  in  composition,  possess  certain  common  properties.  The 
substances,  HC1,  H2S04,  HN03,  H4C202,  and  many  other  compounds 
having  very  different  formulae  are  all  characterized  by  the  possession 
of  a  sour  taste,  the  ability  to  dissolve  zinc  and  other  metals  and  the 
power  of  affecting  certain  organic  dyes  in  a  uniform  way,  as,  for  in- 
stance, turning  blue  litmus  red,  red  phenolphtalein  white  and  yellow 
methyl  orange  red.  Because  of  their  acid  taste,  these  substances  are 
called  acids.  Science  seeks  to  find  a  common  cause  for  a  common 
phenomenon.  The  only  common  factor  present  in  each  of  these  com- 
pounds is  hydrogen. 

On  the  other  hand,  a  great  number  of  organic  compounds  contain 
hydrogen  but  do  not  exhibit  these  properties.  The  hydrogen,  therefore, 
must  be  present  in  two  different  conditions  in  the  two  different  classes  of 
bodies.  In  short,  the  hydrogen  present  in  acids  exists  in  the  ionic 
condition. 

To  the  presence  of  hydrogen  ions  all  acid  properties  are  due,  and 
since  ions  exist  only  in  solution,  it  follows,  strictly  speaking,  that 
acid  properties  are  to  be  found  only  in  solution.  That  this  is  true  is 
indicated  strongly  by  the  facts  that  hydrogen  chlorid,  HC1,  and  hydrogen 
sulfate,  H2S04,  lose  their  acid  properties  so  decidedly  as  they  are  sep- 
arated more  and  more  completely  from  wTater  that  the  conclusion  seems 
evident  that  were  it  possible  to  obtain  these  substances  perfectly  free 
from  water,  they  would  possess  absolutely  no  acid  properties.  The 
addition  of  water,  however,  causes  the  appearance  of  acid  properties  in 
a  marked  degree.  An  acid  may  be  defined  as  any  substance  which  in 
the  presence  of  water  will  dissociate  into  hydrogen  ions. 

The  hydroxids  of  certain  elements  act  in  this  way  and  separate 
hydrogen  ions  in  aqueous  solution.  The  elements  from  which  they  are 
derived  are  known  as  non-metallic  or  acidic,  i.  e.,  acid-forming  elements. 
The  hydrids  of  some  of  these  same  elements  possess  acid  properties. 

The  readiness  with  which  the  various  acids  separate  hydrogen  ions 
and,  consequently,  the  degree  to  which  they  are  present  in  solutions 
varies  greatly  and  constitutes  the  difference  between  strong  and  weak 
acids.  That  acid  is  strongest  which  is  most  completely  dissociated. 


16  QUALITATIVE    ANALYSIS. 

In  solutions  of  such  strength  as  is  suitable  for  qualitative  analysis 
reactions  the  following  acids  are  more  than  three-quarters  dissociated 
into  their  respective  ions.  They  are,  therefore,  very  strong  acids: 

Hydriodic  Chloric  Nitric 

Hydrobromic  Perchloric  Thiocyanic 

Hydrochloric  Permanganic  Tetrathionic 

At  the  same  dilution  the  following  acids  are  about  one-half  dis- 
sociated : 

Sulfuric  Ferrocyamc          Ferricyanic 

Dichromic  Silicofluoric  Oxalic — about  25  per  cent. 

Under  the  same  conditions  the  following  acids  are  dissociated  less 
than  ten  per  cent : 

Sulfurous  Hydrofluoric  Tartaric 

Phosphoric  Nitrous  Acetic 

Arsenic 

None  of  the  following  acids  are  dissociated  as  much  as  one  per  cent. 
in  ordinary  solutions: 

Carbonic  Hydrosulfuric  Boric 

Hypochlorous  Hydrocyanic  Silicic 

Arsenious 

In  dilute  solutions  of  strong  acids  the  dissociation  is  practically 
complete.  When  a  metal  is  added  to  an  acid  in  dilute  solution,  as,  for 
instance,  zinc  and  hydrochloric  acid,  each  H  ion  gives  up  its  plus  charge 
to  the  metal  which  then  passes  into  the  ionic  condition  while  the 
hydrogen  escapes  as  gas,  as  the  following  equation  indicates: 

Zn  +  2H  +  2  Cl  ->  Zn  +  2  Cl  +  H2 

Thus  the  hydrogen  in  an  acid  is  capable  of  being  replaced  by  a 
metal.  There  remains  in  solution  the  same  number  of  Cl  ions  as  before 
the  reaction.  Since  they  appear  on  both  sides  of  the  equation,  they  may 
be  stricken  out  and  the  equation  simplified  as  follows: 


ACIDS.  17 

Thus  it  is  that  hydrogen  is  driven  out  of  the  acid  because  of  the 
greater  tendency  of  the  zinc  to  pass  into  the  ionic  condition.*  In  place 
of  hydrogen  ions  there  exist  in  solution  zinc  ions  which,  on  evaporation 
of  the  solution,  unite  with  the  chlorin  ions  to  form  zinc  chlorid. 

In  hydrochloric  acid,  HC1,  there  is  one  part  of  hydrogen  to  one  of 
chlorin;  in  nitric  acid,  HN03,  one  part  of  hydrogen  to  one  of  the 
group  NO3;  in  sulfuric  acid,  H2S04,  two  parts  of  hydrogen  to  one  of 
the  radicle  S04 ;  in  phosphoric  acid,  H3P04,  three  parts  of  hydrogen  to 
one  of  the  radicle  PO4  and  in  silicic  acid,  H4Si04,  four  parts  of  hydrogen 
to  one  part  of  the  group  Si04.  In  these  compounds,  theoretically  at  least, 
all  of  the  hydrogen  is  capable  of  passing  into  the  ionic  state  and  of 
being  replaced  by  a  metal  in  the  formation  of  salts.  On  the  other  hand, 
in  acetic  acid,  H4C203,  only  one  part  of  hydrogen  is  capable  of  passing 
into  the  ionic  condition ;  in  tartaric  acid,  H6C406,  two  parts  and  in  citric 
acid,  H8C6O7,  only  three  parts.  Hence  it  is  obvious  that  all  of  the  hy- 
drogen in  a  compound  may  not  be  capable  of  passing  into  the  ionic 
condition  and  of  conferring  acid  properties  on  its  solutions. 

Acids  are  classified  according  to  the  amount  of  hydrogen  which 
may  become  ionic.  Eegarding  them  from  the  molecular  standpoint,  if 
from  one  molecule  of  a  substance  one  atom  of  hydrogen  may  pass  into 
the  ionic  condition  and  be  replaced  by  a  metal  in  the  formation  of 
salts,  the  compound  is  said  to  be  a  monobasic  acid.  Dibasic,  tribasic  and 
tetrabasic  acids  separate  two,  three  and  four  hydrogen  ions  from  one 
molecule  of  the  substance  and  are  known  as  polybasic  acids.  It  is  cus- 
tomary to  indicate  that  part  of  the  hydrogen  in  a  compound  which  is 
capable  of  forming  hydrogen  ions.  In  a  formula,  it  is  usually  written 
first,  separated  from  the  rest  of  the  hydrogen  in  such  a  compound,  acetic 
acid  being  written  H(C2H302)  and  tartaric  acid  H2(C4H4O6).  In  the 
dissociation  of  such  acids  that  part  of  the  hydrogen  which  does  not  form 
the  kathion  remains  united  with  the  other  elements  in  a  group  which 
becomes  the  anion. 

All  of  the  hydrogen  in  a  polybasic  acid  does  not  pass  at  the  same 
time  into  the  ionic  condition.  The  dissociation  takes  place  step  by  step 

*  The  varying  tendency  of  elementary  substances  to  take  on  the  electrical  charge 
and  pass  into  the  ionic  condition  is  shown  by  the  table  known  as  the  Potential  Series 
given  on  page  327.  In  this  table  the  elements  are  arranged  in  regular  series  be- 
ginning with  those  which  pass  most  rapidly  into  the  ionic  condition.  As  in  the 
case  of  zinc  and  the  hydrogen  ion,  if  a  free  element  comes  in  contact  with  the  ions 
of  another  element  standing  much  lower  in  the  Potential  Series,  the  charge  passes 
from  the  one  to  the  other  and  the  former  goes  into  the  ionic  condition  while  the 
latter  is  thrown  out  in  the  free  state.  Many  illustrations  of  this  principle  will  be 
noted  in  the  second  part  of  this  book. 

3 


18  QUALITATIVE    ANALYSIS. 

and  the  process  of  separating  one  hydrogen  ion  from  each  molecule  of 
acid  present  is  generally  pretty  well  advanced  before  the  separation  of 
a  second  hydrogen  ion  from  any  molecule  begins.  The  dissociation  of  a 
third  and  fourth  ion  from  the  same  molecule  is  correspondingly  tardy. 
In  dilute  solutions  of  strong  acids,  such  as  the  dibasic  sulfuric  acid,  the 
dissociation  according  to  both  steps  is  practically  complete  and  the  preva- 
lent anion  contains  no  hydrogen,  as  the  following  equations  will  indicate  : 


With  weak  acids  dissociation  of  all  the  hydrogen  present  in  any  one  mole- 
cule remains  to  a  considerable  degree  unaccomplished  even  in  dilute 
solutions.  Thus  the  tribasic  phosphoric  acid  separates  ions  according 
to  the  first  and  second  of  the  following  reactions,  but  the  hydrogen-free 
anion  formed  according  to  the  third  reaction  is  present  in  solution  to 
only  a  small  amount  : 


HP04-»H  +  P04 

Analogous  to  the  acids  there  exists  a  group  of  the  hydroxids  of 
metallic  elements  which  possess  the  common  properties  of  destroying 
acids,  of  turning  litmus  blue  and  of  affecting  other  organic  dyes  uni- 
formly. These  compounds  differ  in  their  composition  and  are  alike  only 
in  that  they  form  hydroxyl  ions,  OH,  in  solution.  These  substances  are 
known  as  bases,  since  by  neutralizing  the  various  acids,  with  each  hy- 
droxid  as  a  base,  a  series  of  compounds,  known  as  salts,  results.  With  the 
bases  as  with  the  acids,  the  degree  of  dissociation  determines  the  strength. 

In  solutions  of  the  strength  ordinarily  used  in  qualitative  analysis, 
sodium  and  potassium  hydroxids  are  more  than  three-quarters  dis- 
sociated, and,  therefore,  are  very  strong  bases. 

Barium  hydroxid  is  about  one-half  dissociated.  Strontium,  calcium 
and  magnesium  hydroxids  are  much  less  soluble,  but  to  the  extent  to 
which  they  dissolve,  they  dissociate  approximately  to  the  same  degree 
as  barium  hydroxid  at  the  same  dilution.  As  bases,  therefore,  they  are 
of  about  the  same  strength. 

Ammonium  and  silver  hydroxid  are  dissociated  less  than  ten  per  cent. 

The  hydroxids  of  the  other  metals  are  very  insoluble  and  act  as  very 
weak  bases,  trivalent  hydroxids  being  generally  weaker  than  divalent 
hydroxids. 


OF 

BASES.  19 

When  an  acid  and  a  base  are  mixed  together  in  the  proper  proportions 
the  one  is  said  to  neutralize  the  other,  for  the  resulting  solution  possesses 
neither  acidic  nor  basic  properties.  On  treating  a  dilute  solution  of 
hydrochloric  acid  with  a  dilute  solution  of  sodium  hydroxid,  such  a 
reaction  takes  place  and  the  acid  is  neutralized  by  the  base.  This  reac- 
tion is  often  written  : 

NaOH  +  HC1  ->  NaCl  +  H20 

In  dilute  solutions,  however,  hydrochloric  acid,  being  a  strong  acid,  and 
sodium  hydroxid,  being  a  strong  base,  are  dissociated  to  practical  com- 
pletion. Sodium  chlorid  is  almost  completely  ionized  at  the  same  dilution. 
Hence  the  reaction  may  be  more  correctly  written  : 

Na  -f  Cl-f  H  +  OH^Na  +  cl+  H20 

After  mixing,  the  sodium  and  chlorin  are  present  in  the  ionic  condition 
exactly  as  they  were  in  the  two  original  solutions.  Since  they  appear 
on  both  sides  of  the  equation,  they  may  be  stricken  out  and  the  equation 
simplified  as  follows: 

H-»HOH 


The  only  product  actually  formed  on  mixing  the  solutions  is  water. 
Since  all  acid  solutions  owe  their  properties  to  the  presence  of  the  hydrogen 
ions  while  basic  solutions  owe  their  properties  to  the  presence  of  hydroxyl 
ions,  the  neutralization  of  (dilute  solutions  of)  strong  acids  with  strong 
bases  is  essentially  the  formation  of  water  from  its  ions. 

To  speak  of  the  formation  of  a  salt  in  this  process  is,  therefore, 
technically  incorrect  since  no  reaction  takes  place  to  form  it,  except  on 
evaporation  of  the  solution  when  the  ions  come  together.  That  this  is 
true  is  indicated  by  the  fact  that  when  equivalent  amounts  of  fully 
dissociated  solutions  of  any  acid  and  base  are  mixed,  the  amount  of  heat 
liberated  is  always  the  same.  When  any  compound  is  formed,  a  definite 
amount  of  heat  enters  the  reaction  which  is  known  as  the  heat  of  forma- 
tion. This  is,  in  general,  different  for  every  compound.  Hence,  if 
different  salts  were  actually  formed  by  mixing  dilute  solutions  of  sul- 
furic,  nitric,  and  hydrochloric  acids  with  dilute  solutions  of  sodium, 
potassium  and  calcium  hydroxids,  the  heats  of  neutralization  could 
not  be  the  same.  On  the  other  hand,  if  one  product  only,  i.  e.,  water,  is 
formed  by  mixing  these  different  reagents,  we  should  expect  the  same 
amount  of  heat  to  be  liberated  in  each  case.  By  the  formation  of  water 
from  its  ions  all  hydrogen  and  hydroxyl  ions  are  removed  and  all  acid 


20  QUALITATIVE   ANALYSIS. 

and  basic  properties  disappear.  There  is  left  a  solution  containing  the 
ions  characteristic  of  a  salt.  Since  by  evaporation  of  the  solution  a  salt 
is  actually  formed,  it  is  ordinarily  stated  that  the  neutralization  of  an 
acid  with  a  base  forms  a  salt. 

Of  the  three  typical  methods  of  salt  formation:  (a)  by  the  action 
of  an  acid  on  a  metal,  (b)  by  the  neutralization  of  an  acid  with  a  base, 
(c)  by  the  action  of  an  acid  on  a  metallic  oxid,  two  have  already  been 
considered.  The  third  method  is  a  process  in  which  the  charges  on  the 
hydrogen  ions  are  transferred  to  a  metallic  kathion  in  a  manner  similar 
to  the  interaction  of  hydrogen  ions  and  metals.  After  passing  out  of 
the  ionic  condition  the  hydrogen  is  not  set  free,  however,  but  unites  with 
the  oxygen  of  the  oxid  to  form  water.  A  typical  reaction  will  make 
clear  the  changes  that  take  place. 

CaO  +  2H->  Ca  +  H20 

It  is  not  essential  that  all  of  the  hydrogen  and  oxygen  contained 
in  a  base  be  capable  of  forming  OH  ions.  As  with  acids,  so  with  bases. 
A  substance  capable  of  forming  one  OH  ion  from  one  molecule  of  the 
compound  is  called  a  monoacid  base  since  in  the  formation  of  salts  this 
OH  ion  is  replaced  by  a  single  anion  of  a  monobasic  acid.  Di,  tri,  and 
tetracid  bases  are  capable  of  separating,  theoretically  at  least,  two, 
three  and  four  OH  ions  from  one  molecule  of  the  compound. 

One  molecule  of  a  monacid  base  will  neutralize  one  molecule  of  a 
monobasic  acid.  Four  molecules  of  a  monobasic  acid  will  react  with  one 
molecule  of  a  tetracid  base.  Two  molecules  of  a  tribasie  acid  will  unite 
with  three  molecules  of  a  diacid  base  and  vice  versa.  Polyacid  bases 
dissociate  step  by  step  in  a  manner  exactly  like  the  ionization  of  poly- 
basic  acids.  (See  page  18.) 

While  there  is  the  greatest  variation  in  the  degree  of  dissociation  of 
the  various  acids  and  bases,  most  salts  ionize  freely.*  Even  the  salts  of 

*  Salts  may  be  divided  according  to  a  classification  based  on  the  acids  and  bases 
from  which  they  are  derived  and  the  variation  in  the  degree  of  dissociation  for  salts 
may  be  stated  for  each  of  these  types. 

Salts  formed  by  the  interaction  of  monobasic  acids  and  monoacidic  bases  dis- 
sociate almost  as  readily  as  the  strongest  acids  and  bases. 

Salts  formed  from  a  dibasic  acid  and  a  monoacidie  base  or  vice  versa  are  some- 
what more  than  half  dissociated  at  ordinary  dilutions. 

Salts  derived  from  dibasic  acids  and  diacid  bases  are  somewhat  less  than  half 
dissociated  under  the  same  conditions. 

With  salts  derived  from  tri-  and  tetrabasic  acids  and  tri-  and  tetra-acidic  bases 
other  phenomena  often  play  more  important  parts  than  dissociation.  (See  pages 
28-30,  also  46.) 


SALTS.  21 

weak  acids  and  weak  bases  dissociate  readily.  In  this  case,  however, 
other  factors  enter  into  the  consideration  and  complicate  the  result. 
(See  pages  28-30.) 

Acids  may  be  designated  as  those  substances  which  separate  hydrogen 
as  kathion.  Bases  separate  hydroxyl  as  anion.  Salts  separate  kathions 
other  than  hydrogen  and  anions  other  than  hydroxyl.  Many  substances 
are  known,  however,  which,  in  solution,  separate  both  hydroxyl  and 
other  non-metallic  anions  at  the  same  time.  These  are  evidently  mix- 
tures of  the  ions  characteristic  of  bases  and  of  salts.  In  some  instances 
possibly  no  serious  objection  can  be  raised  to  considering  such  sub- 
stances as  very  intimate  mixtures  of  the  base  with  one  of  its  salts.  In 
other  cases  the  mixture  seems  to  be  inside  the  molecule.  Each  individual 
molecule  seems  to  possess  the  properties  of  both  base  and  salt.  The 
substances  act  as  though  they  were  in  no  ordinary  sense  mixtures,  since 
they  cannot  by  any  means  be  separated  into  the  pure  salt  and  the  pure 
base.  On  further  treatment  with  acids  they  lose  their  basic  properties 
entirely.  For  these  reasons  they  are  believed  to  be  salts  in  which  part 
only  of  the  hydroxyl  which  is  characteristic  of  the  base  has  been 
replaced  by  an  acid  anion.  They  are  called  basic  salts,  therefore,  and 
the  formulae  indicate  their  apparent  constitution. 

OH 

/ 

—  OH  Pb 

Fe  — OH  \ 

-  (C2H302)  (C2H302) 

Basic  ferric  acetate.  Basic  lead  acetate. 

Analogously,  certain  other  substances  separate  at  the  same  time 
hydrogen  and  other  metallic  kathions.  For  similar  reasons  they  are 
known  as  acid  salts.  Since  the  proportion  of  salt  to  the  acid  is  far  more 
constant  than  the  proportion  of  salt  to  base  in  basic  salts,  there  hardly 
seems  reason  for  considering  any  well-characterized  acid  salt  as  a  mere 
mixture  of  an  acid  with  one  of  its  salts.  The  following  formulae  indicate 
the  apparent  constitution  of  typical  acid  salts. 

Certain  exceptions  to  this  general  statement  are  well  known.  The  halogen  com- 
pounds of  mercury  ionize  to  a  very  slight  degree,  those  of  cadmium  somewhat  more 
while  the  zinc  halids  together  with  lead  and  certain  other  acetates  stand  inter- 
mediate between  these  and  other  salts. 


22  QUALITATIVE    ANALYSIS. 

H 

C03 

K  Na  Ca 

\  \  \ 

S04  C03  C03 

/  /  / 

H  H  H 

Acid  potassium  sulfate.       Acid  sodium  carbonate.       Acid  calcium  carbonate. 

Acid  salts  are  frequently  indicated  by  the  prefix  "bi,"  as  potassium 
bisulfate,  sodium  bicarbonate,  etc. 

To  distinguish  acid  and  basic  salts  from  the  type  which  separates 
only  kathions  other  than  hydrogen  and  anions  other  than  hydroxyl, 
members  of  this  latter  class  are  called  normal  salts. 

Acid  salts  are  usually  far  more  soluble,  basic  salts  far  less  soluble, 
than  the  corresponding  normal  salts. 

Certain  bodies  are  capable  of  dissociating  in  more  than  one  way. 
The  hydroxids  of  elements  which  have  neither  a  pronounced  metallic 
nor  non-metallic  character  may  separate  under  one  condition  hydrogen 
ions  and  under  another  condition  hydroxyl  ions.  Arsenic,  antimony, 
tin,  lead,  zinc  and  aluminum  are  elements  of  this  nature  and  their 
hydroxids  may  act  either  as  acids  or  bases.  To  illustrate:  zinc  hy- 
droxid,  Zn02H2,  may  dissociate  into  Zn  and  OH  ions  in  which  case 
the  positively  charged  zinc  ions  are  acting  as  metallic  ions.  Zinc 
hydroxid  may  separate  H  ions  and  Zn02  ions  in  which  case  the  nega- 
tively charged  anion  shows  zinc  to  be  playing  the  part  of  an  acid- 
forming  element.  In  acid  solutions  zinc  hydroxid  separates  OH  ions 
which  unite  with  the  H  ions  furnished  by  the  acid  to  form  undissociated 
water  and  leave  in  solution  ions  of  a  salt,  such  as  zinc  chlorid,  ZnCl2. 
In  solutions  containing  hydroxyl  ions  in  quantity,  zinc  hydroxid  sepa- 
rates H  ions.  The  hydrogen  and  hydroxyl  ions  unite  to  form  water  and 
leave  in  solution  sodium  zincate,  Na2Zn02,  or  other  salt  in  which  zinc 
acts  as  an  acid-forming  element. 

EQUILIBRIUM   AND    MASS   ACTION. 

If  the  metathesis  AX  +  BY  <-»  AY  +  BX  be  a  reversible  reaction, 
when  AX  and  BY  are  brought  together,  they  will  react  to  form  AY 
and  BX  with  a  definite  reaction  speed.  At  the  end  of  the  first  short 
period  during  which  the  reaction  has  run  there  will  exist  an  intimate 
mixture  of  four  substances  of  which  AX  and  BY  will  be  present  in 


EQUILIBEIUM.  23 

much  greater  proportion  than  AY  and  BX.  AY  and  BX-  as  soon  as 
formed  will  react  together  to  form  AX  and  BY  again,  but,  other  things 
being  equal,  more  of  AY  and  BX  will  be  formed  in  any  given  interval 
by  the  original  metathesis  than  of  AX  and  BY  by  the  second  reaction, 
because  of  the  simple  fact  that  there  is  so  much  more  of  AX  and  BY 
to  interact  than  there  is  of  AY  and  BX.  The  first  reaction  will  continue 
to  predominate  until  all  four  substances  are  present  in  certain  definite 
proportions.  Then  we  may  consider  each  reaction  as  taking  place  at  the 
same  time  and  to  such  a  degree  that  the  proportions  of  the  constituents 
remain  unchanged. 

If,  now,  more  of  AX  and  BY  be  added,  the  original  reaction  will 
again  become  dominant  until  the  same  proportion  among  the  con- 
stituents is  reached. 

Such  a  condition  which)  when  disturbed,  tends  to  restore  itself  is  known 
as  equilibrium,  and  the  reaction  by  which  it  is  restored  as  the  equi- 
librium reaction. 

Such  is  the  nature  of  the  process  which  is  going  on  whenever  a 
chemical  reaction  takes  place.  That  reactions  are  reversible  generally 
is  believed  to  be  true.  The  fact  that  many  reactions  do  not  seem  to  be 
reversible  in  practice  is  due  to  the  presence  of  determining  factors  which 
shift  the  equilibrium,  from  the  condition  of,  possibly,  equal  proportions 
among  the  constituents  to  a  condition  in  which  some  constituents  are 
practically  absent. 

To  consider  a  particular  reaction :  If  ethyl  hydroxid  (i.  e.,  ordinary 
alcohol)  and  acetic  acid  be  heated  together,  a  reaction  takes  place  in 
part,  leading  to  the  formation  of  ethyl  acetate  and  water.  Equilibrium 
is  soon  reached  in  which  the  four  substances,  alcohol,  acetic  acid,  ethyl 
acetate  and  water  are  present  in  proportions  which  are  influenced  by  tem- 
perature, but  which  are  definite  for  any  given  temperature.  On  the 
other  hand,  if  ethyl  acetate  and  water  are  heated  together,  a  partial  re- 
action ensues  leading  to  the  formation  of  some  alcohol  and  acetic  acid. 
Equilibrium  ensues,  a  mixture  of  the  four  substances  being  formed 
which  has  the  same  proportionate  composition  as  in  the  first  reaction, 
provided  the  temperature  be  the  same.  Thus  it  makes  no  difference  from 
which  end  the  start  is  made,  a  condition  of  equilibrium  results  in  which 
the  amounts,  or  concentrations,  of  the  components  are  the  same. 

These  relations  have  been  proved  to  be  so  constant  that  they  have 
found  mathematical  expression.  The  product  of  the  concentrations  of 
all  the  factors  on  one  side  of  an  equation  is  equal  to  the  product  of  the 
concentrations  of  all  the  factors  on  the  other  side  of  the  equation  multi- 


24  QUALITATIVE    ANALYSIS. 

plied  ~by  a  constant  quantity.     This  may  be  expressed  algebraically  as 
follows 

A  X  B  X  C  X  D   ect.  =X  X  Y  X  Z   ect.  .  .  .  X  K 
or 

A  X  B  X  C  ect. 


X  X  Y  X  Z  ect. 


T7" 


where  A,  B,  C,  X,  Y,  Z  ect.,  refer  to  the  concentrations  of  the  different 
substances,*  molecules,  atoms  or  ions  as  the  case  may  be,  which  enter  into 
the  reaction,  and  K  is  a  constant  quantity  depending  on  conditions  for 
its  value  but,  once  determined,  remaining  unchanged  for  these  condi- 
tions. This  equation  is  known  as  the  equilibrium  equation. 

Let  the  condition  of  equilibrium  in  the  reaction  already  considered 
be  expressed  by  the  equation 

AXB=XXYXK 

in  which  A  represents  the  concentration  of  alcohol,  B  of  acetic  acid,  X 
of  ethyl  acetate  and  Y  of  water.  If  one  of  these  factors,  as  A,  be  in- 
creased, the  product  AB  will  be  larger  than  the  product  XYK  and  the 
algebraic  expression  will  not  be  true  without  a  readjustment  of  values 
on  both  sides. 

Just  as  the  algebraic  equation  is  rendered  untrue  by  the  increase  in 
one  of  its  factors,  so  is  the  chemical  equilibrium  disturbed  by  the  increase 
in  the  concentration  of  one  of  its  components.  A  reaction  will  take  place 
by  virtue  of  which  the  concentration  of  the  factors  A  and  B  will  be 
lessened  while  the  concentration  of  X  and  Y  will  be  increased  until 
equilibrium  ensues,  when  the  equation  will  be  true  again. 

These  relations  hold  in  practice.  If  it  is  desirable  to  convert  one 
gram  of  acetic  acid  entirely  into  ethyl  acetate,  it  has  been  found 
advantageous  to  increase  decidedly  the  amount  of  alcohol  present 
beyond  that  which  is  sufficient  to  unite  with  the  acid.  Because  of  the 
mass  of  the  alcohol  present  it  causes  the  reaction  to  run  in  the  direction 
which  will  tend  to  reduce  its  concentration,  i.  e.,  to  use  it  up  by  forming 
ethyl  acetate.  Thus  is  equilibrium  between  the  components  brought 
about. 

*  The  equilibrium  equation  must  contain  a  factor  representing  each  unit  (mole- 
cule, atom  or  ion)  which  enters  into  the  reaction  whether  they  be  the  same  or  dif- 
ferent. Hence  if  two  molecules  (atoms  or  ions)  of  the  same  substance  appear  in 
the  equation,  the  factor  representing  the  concentration  of  this  substance  must  be 
repeated  as  A  X  A,  i.  e.,  A2.  If  three  molecules  (atoms  or  ions)  of  the  same  sub- 
stance enter  the  reaction,  the  factor  becomes  A3. 


MASS    ACTION.  25 

Such  action  of  a  reagent  in  shifting  the  condition  of  equilibrium 
because  of  the  large  excess  in  which  it  is  present  is  known  as  mass  action. 
By  increasing  the  concentration  of  the  components  on  one  side  of  an 
equation,  or  conversely,  by  decreasing  the  concentration  of  the  com- 
ponents on  the  opposite  side,  mass  action  may  be  brought  about. 

Considering  the  same  reaction  after  equilibrium  has  been  reached, 
if  the  concentration  of  water  present  be  decreased,  the  equilibrium  will 
be  disturbed  and  the  alcohol  and  acid  will  combine.  This  will  lessen  the 
concentrations  on  the  side  of  the  larger  product  and  increase  the  con- 
centrations on  the  smaller  side,  thus  restoring  equilibrium. 

To  consider  another  very  instructive  illustration:  The  process  of 
dissociation  is  a  typical  reversible  reaction  which  always  comes  to  a 
condition  of  equilibrium  between  the  ionized  and  un-ionized  portions  of 
the  substance  in  solution.  The  degree  of  dissociation,  i.  e.,  the  ratio 
between  the  ionized  and  non-ionized  portions,  depends  largely  upon  the 
dilution  of  the  solution.  This  is  a  direct  consequence  of  the  mass  action 
relation,  as  can  be  demonstrated  readily. 

Let  common  salt  be  put  into  aqueous  solution  to  an  amount  sufficient 
to  make  the  concentration  of  the  dissolved  substance  equal  to  unity. 
Concentrations  are  expressed  in  molecules,  atoms  or  ions  of  the  active 
agents  in  a  reaction  as  the  case  may  be.  Whenever  the  same  number  of 
molecules,  atoms  or  ions  are  present  in  a  given  volume,  the  concentrations 
are  equal.  Suppose  the  dilution  to  be  such  that  the  salt  is  one-half  dis- 
sociated. The  number  of  undissociated  molecules  is  one-half  the  total 
number  of  molecules  added;  the  concentration  of  undissociated  salt  is, 
therefore,  J.  The  number  of  molecules  of  the  dissociated  salt  is  equal 
to  the  number  of  molecules  undissociated.  But  each  molecule  of  NaCl 

+ 

that  dissociates  separates  one  Na  and  one  Cl  ion.  Therefore,  the  num- 
ber of  Na  ions  is  equal  to  the  number  of  undissociated  molecules;  the 
concentration  of  Na  ions  is,  therefore,  -|.  Chlorin  ions  are  present  to 
the  same  concentration  also.  The  general  equation  for  the  dissociation  of 
a  binary  electrolyte,  i.  e.,  one  which  forms  two  ions  only,  is 

A=XXYXK 

It  becomes  in  this  instance 

i  =  JXiXK 
in  which  case  K  =  2. 

Suppose,  now,  that  the  solution  be  diluted  to  twice  its  volume.  The 
number  of  molecules  of  undissociated  salt  in  unit  volume  of  solution  is 
now  one-half  of  its  former  value.  Its  concentration  will  be  one-half  the 
former  concentration,  or  J.  Similarly  the  concentrations  of  Na  and  Cl 


26  QUALITATIVE    ANALYSIS. 

ions  will  be  halved.     Since  the  constant  quantity  K  is  not  influenced  by 
dilution  of  the  solution,  it  will  be  unchanged.    The  equation  becomes  then  : 


which  is  untrue.  The  product  of  the  factors  on  the  right  side  is  much 
smaller  than  the  concentration  of  the  left  side.  lonization  takes  place, 
reducing  the  concentration  of  undissociated  salt  and  increasing  the  con- 
centrations of  the  ions  until  the  equation  is  once  more  true.  Thus 
dissociation  is  directly  dependent  on  the  dilution  of  the  solution. 

When  the  volume  of  the  solution  is  so  large  that  further  dilution 
gives  no  evidence  of  further  dissociation,  the  substance  is  said  to  be  in 
a  state  of  complete  dissociation.  According  to  these  relations,  however, 
there  must  still  be  present  a  minute  amount  of  undissociated  salt,  the 
concentration  of  which  is  practically  zero. 

The  conceptions  of  equilibrium  and  mass  action  lie  at  the  foundation 
of  any  understanding  of  chemical  reactions.  Operating  according  to  the 
relations  just  explained,  certain  physical  properties  of  components  in 
the  reactions  often  act  as  determining  factors  and  cause  the  condition 
of  equilibrium  to  be  such  that  the  concentration  of  certain  factors  be- 
comes practically,  if  not  actually,  zero.  Because  of  this  relation  we  can 
predict  that  whenever  a  gaseous,  insoluble  or  slightly  dissociating  sub- 
stance can  be  formed,  it  will  be  formed  so  long  as  there  are  factors  present 
capable  of  forming  it.  This  statement  applies  chiefly  to  reactions  taking 
place  in  solution  but  the  formation  of  a  volatile  product  is  a  determining 
factor  in  reactions  performed  in  the  "dry  way"  as  well. 

Let  us  consider  the  substances  AX  and  BY  in  solutions  so  dilute  as 

+  + 

to  be  completely  dissociated  into  the  kathions  A  and  B  and  the  anions 

X  and  Y.    On  partial  evaporation  of  the  mixed  solutions,  in  the  absence 

+ 

of  any  determining  factors,  the  kathion  A  would  unite  with  the  anions 
X  and  Y  in  proportions  determined  by  the  concentrations  of  the  various 
ions  present.  In  a  similar  way  B  will  unite  with  X  and  Y,  giving  a 
mixture  of  AX,  AY,  BX  and  BY.  A  condition  of  equilibrium  will 
result  in  which  these  four  substances  may  be  considered  as  ionizing  and 
reacting  together  in  all  possible  ways,  but  to  such  a  degree  that  the 
proportions  between  the  constituents  are  not  changed. 

If,  however,  one  of  the  substances,  as  AY,  be  gaseous,  it  will,  as 
soon  as  formed,  pass  out  of  the  solution  and  be  incapable  of  affecting 
the  reaction  in  any  way.  By  the  fact  that  AY  is  volatile  its  concentra- 
tion is  kept  continually  reduced  to  practical  zero.  The  ions  A  and  Y 
keep  coming  together  to  form  more  undissociated  salt  in  the  endeavor 


DETEEMINING   FACTOES.  27 

to  bring  about  an  equilibrium  which  is  continually  disturbed  by  the  vola- 
tility of  AY.  If  the  reaction  speed  is  sufficiently  great,  all  of  AY  will 
be  formed  in  a  very  short  interval  of  time  and  will  pass  out  of  solution. 
There  will  be  left  in  solution  'nothing  but  B  and  X  to  form  undissociated 
BX  on  evaporation  of  the  solution.  Thus  a  gaseous  product  determines 
the  direction  in  which  a  reaction  will  go  by  keeping  the  concentration  of 
the  constituents  forming  it  reduced  to  zero. 

An  insoluble  substance  acts  in  exactly  the  same  way.  If  the  reac- 
tion be  one  that  takes  place  in  solution  only,  an  insoluble  substance 
suspended  in  the  liquid,  however  finely  divided,  will  no  more  be  present 
in  solution  or  affect  the  reaction  than  a  glass  rod  in  the  beaker.  The 
difference  between  the  insoluble  substance  inside  the  beaker  and  sur- 
rounded by  solution  and  the  gaseous  substance  unquestionably  removed 
from  the  sphere  of  action  is  apparent  and  not  real. 

If  AY  be  a  non-volatile  soluble  substance  dissociating  but  little 
in  comparison  with  AX,  BX  and  BY,  it  determines  the  reaction  none 
the  less  definitely.  Most  reactions  in  solution  take  place  between  ions 

only.     If  AY  does  not  dissociate  appreciably,  the  concentrations  of  ions 

+ 

A  and  Y  is  kept  less  than  is  necessary  for  equilibrium  and  they  are 
removed  from  the  sphere  of  action  just  as  completely  as  if  AY  were 
gaseous  or  insoluble. 

Certain  of  these  determining  factors  may  exercise  an  influence  under 
one  condition  in  which  a  reaction  takes  place  and  not  under  other 
conditions.  Hence  the  products  of  a  reaction  may  be  very  different  under 
one  condition  than  under  another,  although,  the  matter  present  may  be 
the  same.  This  point  is  illustrated  on  a  world  scale  by  the  fact  that 
the  characteristic  compounds  of  igneous  rocks  are  anhydrous  silicates 
while  hydrated  silicates  and  carbonates  are  typical  of  all  later  forma- 
tions. This  difference  is  due  to  the  fact  that  at  high  temperatures  water 
and  carbon  dioxid  are  both  volatile  and  were  thus  removed  from  the 
sphere  of  action.  Under  the  conditions  maintaining  on  the  earth  to-day 
water  is  an  ever-present  factor.  Carbon  dioxid  is  quite  soluble  in  it 
and  therefore  the  process  of  hydration  and  decomposition  of  the  silicates 
known  as  "weathering"  is  accompanied  by  the  formation  of  carbonate 
deposits.  With  the  temperature  the  equilibrium  shifts.  That  which 
under  former  conditions  was  the  stable  phase  is  to-day  unstable. 

In  a  similar  way,  from  boric  acid  and  sodium  carbonate  in  solution 
sodium  tetraborate,  Na2B407.10H2O,  crystallizes,  while  at  the  fusion 
temperature  sodium  metaborate,  NaB02,  is  formed. 

In  the  present  consideration  it  has  been  assumed  that  the  solutions 


28  QUALITATIVE    ANALYSIS. 

of  the  reagents  AX  and  BY  are  so  dilute  that  the  dissociation  is  com- 
plete. This  is  by  no  means  essential  to  the  proposition.  In  concentrated 
solutions  a  small  fraction  only  exists  in  the  ionized  condition ;  the  propo- 
sition is  true,  nevertheless.  As  the  concentration  of  the  ions  A  and  Y 
are  reduced  to  zero  by  the  volatility,  insolubility  or  non-dissociation  of 
AY,  the  equilibrium  between  the  dissociated  and  undissociated  portions 
of  AX  and  BY  is  disturbed  in  each  case.  This  leads  to  further  dissocia- 
tion of  each  salt  until  finally  all  of  AX  and  BY  have  been  ionized  and 

reacted.     By  the  operation  of  the  same  principles  in  the  reverse  direc- 

+ 

tion  the  ions  B  and  X  which  are  left  in  solution  may  become  undis- 
sociated BX  to  a  large  degree. 

The  actual  number  of  ions  present  in  solution  at  any  one  time  may  be 
but  a  small  fraction  of  the  total  number  of  ions  that  might  exist  at  a 
greater  dilution.  It  is  not  the  actual  number  of  ions  present  that  de- 
termines the  extent  to  which  a  reaction  will  take  place  or  the  amount  of 
product  that  will  be  formed,  but  rather  the  total  number  of  ions  of  a 
given  kind  that  it  is  possible  to  form  during  the  course  of  the  reaction. 
Ions  not  actually  existent  at  any  one  time  but  possible  of  formation  dur- 
ing the  course  of  a  reaction  are  known  as  potential  ions. 

In  practical  work  it  is  not  sufficient  ordinarily  to  rely  on  reactions 
in  which  part  only  of  the  reagents  interact  while  an  appreciable  fraction 
is  unchanged.  Reactions  of  this  kind,  moreover,  do  not  generally  give 
definite  visible  evidence  that  a  reaction  is  taking  place.  For  use  in 
analytical  chemistry  such  reactions  are  chosen  as  contain  within  them- 
selves some  determining  factor  to  cause  the  practical  completion  of  the 
reaction  and  at  the  same  time  give  visible  evidence  that  a  change  has 
taken  place.  Hence,  such  reactions  as  are  accompanied  by  the  formation 
of  a  precipitate,  the  evolution  of  a  gas  or  a  change  in  color  are  primarily 
important  in  formulating  a  scheme  of  analysis. 

HYDROLYSIS. 

To  learn  something  of  the  application  of  the  principles  of  equilibrium 
and  mass  action  to  the  phenomena  of  dissociation,  let  us  consider  the 
ionization  of  what  may  be  termed  unbalanced  salts.  By  this  term  is 
meant  normal  salts  formed  by  the  action  of  strong  bases  on  weak  acids 
or  the  reverse.  Such  salts  are  to  be  distinguished  from  the  balanced 
normal  salts  formed  by  the  neutralization  of  acids  and  bases  which  dis- 
sociate to  approximately  the  same  degree.  Unbalanced  salts  do  not 
contain  within  themselves  hydrogen  or  oxygen  capable  of  forming  either 
hydrogen  or  hydroxyl  ions  and  they  are,  therefore,  entirely  different 


HYDROLYSIS.  29 

from  basic  or  acid  salts.  Sodium  carbonate  is  such  an  unbalanced  salt, 
being  formed  by  the  action  of  the  strong  base,  sodium  hydroxid,  on  the 
weak  carbonic  acid. 

Let  it  be  remembered  that  a  strong  base  or  acid  dissociates  with 
great  readiness  while  a  weak  base  or  acid  dissociates  but  little.  Salts 
ionize  readily,  however.  Sodium  carbonate,  Na2C03,  may  be  taken  as  a 
type  of  all  other  salts  of  strong  bases  with  weak  acids.  When  it  goes 
into  solution,  it  ionizes  to  a  great  extent,  giving  at  once  a  high  concentra- 
tion of  sodium  and  carbonate  ions. 

The  neutralization  of  dilute  acid  with  dilute  solutions  of  the  strong 
bases  has  been  set  forth  as  consisting  essentially  of  the  union  of  hydroxyl 
ions  with  hydrogen  ions  to  form  water.  The  formation  of  water  from  its 
ions  of  course  implies  dissociation  into  these  ions.  This  is  a  fact  and 
water  may  be  considered  either  as  a  very  weak  acid  or  base.  Water  is  to 
be  remembered,  however,  as  a  substance  which  dissociates  to  an  extremely 
small  degree,  only  about  one  or  two  in  a  thousand  million  molecules  being 
ionized.  Compared  with  compounds  which  dissociate  readily  this  be- 
comes relatively  zero  and  may  be  entirely  disregarded.  In  dealing  with 
substances  which  dissociate  but  little,  these  values  often  become  signifi- 
cant. It  is  so  in  the  present  consideration. 

Sodium  carbonate  introduces  into  solution  a  high  concentration  of 
C03  ions  and  of  Na  ions.  Water  supplies  a  relatively  small  concentra- 
tion of  OH  and  H  ions.  With  these  four  ions  as  factors,  there  is  the 
possibility  of  the  formation  of  NaOH  and  H2C03.  Sodium  hydroxid 
ionizes  readily ;  therefore  the  undissociated  substance  will  not  be  formed, 
but  the  ions  will  stay  in  solution  in  undiminished  concentration.  -  Car- 
bonic acid,  on  the  other  hand,  dissociates  to  an  exceedingly  small  degree. 
The  product  of  the  large  factor  —  the  concentration  of  C03  ions  —  by 
the  small  factor  —  the  concentration  of  H  ions  —  is  too  great  for  the  equi- 
librium between  the  dissociated  and  undissociated  H2C03.  A  reaction 
takes  place  to  decrease  the  concentrations  of  the  ions  and  to  increase 
the  concentration  of  the  undissociated  substance.  The  ions  come  to- 
gether to  form  undissociated  carbonic  acid. 

By  this  removal  of  hydrogen  ions  from  the  sphere  of  action,  the 
equilibrium  between  the  dissociated  and  undissociated  water  is  dis- 
turbed and  more  water  dissociates.  Thus  the  concentration  of  H  ions 
is  raised  to  quite  the  original  figure.  This  fact  again  causes  the  formation 
of  more  undissociated  H2CO3.  The  dissociation  of  water  and  the  forma- 
tion of  undissociated  acid  from  its  ions  continues  until  equilibrium  is 


30  QUALITATIVE    ANALYSIS. 

reached.  There  will  remain  an  excess  of  OH  ions  which  gives  the  solu- 
tion the  characteristic  basic  reaction  to  litmus  and  other  dyes. 

When  a  salt  is  thus  broken  down  by  means  of  water  into  the  cor- 
responding acid  and  base,  the  process  is  called  hydrolysis.  Hydrolysis 
is  a  general  phenomenon  taking  place  whenever  a  salt  of  a  weak  acid  or 
base  goes  into  solution.  Under  ordinary  conditions,  because  of  the 
exceedingly  small  amount  of  dissociated  water  present,  the  proportion 
of  salt  actually  hydrolyzed  is  very  small,  ordinarily  not  over  one  per 
cent.  This  proportion  is  sufficient  to  be  readily  indicated  by  litmus 
paper,  however.  In  some  cases  the  presence  of  the  free  acid  may  be 
indicated  by  the  odor,  although  the  solution  gives  the  basic  reaction. 
Solutions  of  the  alkali  salts  of  the  volatile  hydrocyanic,  hydrosulfuric 
and  sulfurous  acids  exhibit  these  phenomena  to  a  considerable  degree. 

If  solutions  of  the  salts  of  weak  bases  with  strong  acids  be  investi- 
gated, exactly  the  reversed  phenomena  will  be  observed.  Ferric  chlorid 
is  such  a  salt,  ferric  hydroxid  being  a  very  weak  base  while  hydrochloric 
acid  dissociates  readily.  On  going  into  solution  ferric  chlorid  gives  high 

++4- 

concentrations  of  Fe  and  Cl  ions.  The  dissociated  water  gives  low  con- 
centrations of  OH  and  H  ions.  The  product  of  the  large  factor  —  the 
concentration  of  ferric  ions  —  by  the  small  factor  —  the  concentration 
of  OH  ions  —  is  too  great  for  the  equilibrium  between  dissociated  and 
undissociated  ferric  hydroxid.  The  equilibrium  reaction  takes  place, 
reducing  the  concentration  of  the  dissociated  salt  by  the  formation  of  the 
undissociated  base,  Fe(OH)3.  An  excess  of  H  ions  remains  that  gives 
to  the  solution  an  acid  reaction. 

The  fact  that  ferric  hydroxid  is  insoluble  aids  this  reaction  very 
materially.  Ferric  hydroxid  accumulates  as  a  precipitate,  therefore, 
until  the  concentration  of  acid  becomes  sufficient  to  dissolve  it  and  form 
ferric  chlorid  and  water  once  more.  Thus  the  system  comes  to  equi- 
librium. It  is  maintained  in  this  condition  by  the  continuous  formation 
of  free  acid  and  base  by  hydrolysis,  and  the  formation  of  the  salt  by  the 
neutralization  of  the  one  by  the  other.  It  is,  however,  only  necessary 
to  remove  the  acid  by  appropriate  means  in  order  that  complete  hydrol- 
ysis may  take  place  and  all  the  ferric  hydroxid  be  precipitated. 

//  there  be  a  considerable  difference  between  the  strengths  of  the  acid 
and  the  basic  constituent  of  a  salt,  hydrolysis  will  give  to  its  aqueous 
solution  the  reaction  of  the  constituent  which  dissociates  to  the  greater 
extent. 


KEPEESSION    OF    IONIZATION.  31 

REPRESSION    OF    IONIZATION. 

In  dilute  solutions  the  dissociation  of  a  substance  is  not  ordinarily 
affected  appreciably  by  other  substances  present  in  reasonably  small 
amounts,  provided  they  do  not  combine  directly  with  any  of  the  ions  pres- 
ent. If,  however,  two  substances  are  capable  of  forming  a  common  ion, 
they  will  have  a  decided  effect  on  each  other  especially  if  one  dis- 
sociates more  readily  than  the  other.  Acetic  acid  and  its  salts  illustrate 
this  phenomenon  very  satisfactorily  and  furnish  an  instance  of  great 
interest  to  analytical  chemistry. 

In  dilute  acetic  acid  solutions  only  small  concentrations  of  hy- 
drogen and  acetate  ions  are  present,  since  at  the  ordinary  dilution,  the 
ionization  of  acetic  acid  is  only  four  per  cent,  of  that  given  by  hydro- 
chloric acid.  For  this  reason  many  reactions  will  take  place  in  acetic 
acid  solutions  which  cannot  be  brought  about  in  solutions  of  the  mineral 
acids.  In  the  equation  representing  the  equilibrium  between  the  dis- 
sociated and  undissociated  acid,  the  factors  representing  the  concentra- 
tions of  the  hydrogen  and  acetate  ions  are  equal. 

If  a  considerable  amount  of  sodium  acetate  or  other  salt  of  acetic 
acid  be  added  to  the  solution,  since  salts  dissociate  readily,  a  high  con- 
centration of  metallic  ions  and  of  acetate  ions  is  introduced.  The  total 
concentration  of  acetate  ions  is  now  many  times  the  concentration  pro- 
duced when  the  acetic  acid  first  came  to  equilibrium  with  its  ions. 
Consequently,  the  initial  equilibrium  is  disturbed  and  a  reaction  ensues 
to  lessen  the  concentration  of  the  ionized  portion  of  the  acid.  Hydrogen 
and  acetate  ions  come  together,  an  equal  number  of  each,  to  form  undis- 
sociated acid.  But  since  the  concentration  of  acetate  ions  was  increased 
by  a  very  considerable  factor,  the  concentration  of  hydrogen  ions  must 
be  decreased  correspondingly  until  equilibrium  is  reached. 

In  the  equation  representing  this  final  condition  of  equilibrium  the 
concentration  of  acetate  ions  is  a  very  large  factor  while  the  concentra- 
tion of  hydrogen  ions  is  very  small.  In  this  way  the  concentration 
of  hydrogen  ions,  consequently  the  acidity  of  the  solution,  is  greatly 
reduced  without  neutralization  or  removal  of  the  acid.  This  phenomenon 
is  known  as  the  repression  of  ionization.  Whenever  a  strongly  disso- 
ciating electrolyte  is  added  to  a  solution  of  a  slightly  dissociating  sub- 
stance, the  ionization  of  the  weaker  one  is  thereby  repressed,  provided 
the  two  substances  have  a  common  ion. 

As  a  further  illustration  of  this  same  principle  the  action  of  salts  of 
ammonium  upon  the  free  base  may  be  cited.  Ammonium  hydroxid  is 
a  weak  base,  at  ordinary  dilutions  dissociating  only  to  about  one-four- 


32  QUALITATIVE    ANALYSIS. 

hundredth  of  the  extent  to  which  sodium  and  potassium  hydroxids  ionize. 
By  the  addition  of  the  chlorid  or  other  salt  of  ammonium  to  its  solutions 
the  concentration  of  hydroxyl  ions  may  be  much  further  reduced.  Am- 
monium chlorid  dissociates  readily,  introducing  a  high  concentration  of 
ammonium  ions  which  disturbs  the  original  equilibrium  between  the 
undissociated  base  and  its  ions.  Ammonium  ions  unite  with  hydroxyl 
ions,  therefore,  to  increase  the  concentration  of  the  undissociated  base 
and  decrease  the  concentration  of  the  ions.  By  the  removal  of  hydroxyl 
ions  from  the  solution  its  basic  properties  are  lessened.  This  principle 
finds  frequent  application  in  the  separation  of  the  different  metals  in 
qualitative  analysis.  (See  pages  265  and  269.) 

SOLUTE   AND    SOLVENT. 

When  a  substance,  technically  known  as  the  solute,  is  brought  in 
contact  with  a  solvent,  it  dissolves,  the  proportion  of  solute  to  solvent 
increases  from  zero*  until  a  certain  definite  ratio  between  the  solute  and 
the  solvent  is  reached  which  is  known  as  the  solubility  of  the  substance 
dissolved.  The  solubility  of  a  substance  expresses  the  proportionate 
composition  of  a  solution  which  is  at  equilibrium  with  the  solute  in  con- 
tact with  it;  that  point  at  which  as  much  goes  into  solution  as  separates 
from  it  in  any  given  interval  of  time.  A  solution  which  is  at  equilibrium 
with  the  solute  is  said  to  be  a  saturated  solution. 

Most  solids  dissolve  to  a  greater  degree  in  hot  solvents  than  in  cold 
solvents  and  the  solubility  is  said  to  increase  with  the  temperature. 
Liquid  solutes  are  affected  much  less  regularly  by  temperature,  while 
the  solubility  of  gaseous  solutes  decreases  as  the  temperature  increases. 

If  a  solution  of  a  solid  prepared  at  a  high  temperature  be  cooled, 
as  its  solubility  decreases  with  the  temperature,  it  will  contain  more  of 
the  solute  than  it  can  hold  and  still  be  in  equilibrium.  A  solution  in 
this  condition  is  said  to  be  a  supersaturated  solution. 

The  normal  tendency  of  such  a  supersaturated  solution  is  to  sep- 
arate the  excess  from  solution.  This  readily  occurs  if  some  of  the 
solid  be  in  contact  with  the  solution.  If  no  solid  solute  be  present, 
however,  solutions  of  some  substances  show  a  surprising  tendency  to  re- 

*  As  the  amount  of  solvent  contained  in  a  given  amount  of  solute  increases  from 
zero  the  specific  gravity  of  the  solution  changes  slowly  and  may  become  either  greater 
or  less  than  the  specific  gravity  of  the  pure  solvent.  At  any  given  temperature  there 
is  a  definite  specific  gravity  for  each  definite  proportion  between  solute  and  solvent. 
Many  common  substances  have  been  carefully  investigated  in  this  respect.  Hence 
it  is  possible  to  ascertain  the  concentration  of  a  solution  of  these  substances  simply 
by  determining  the  specific  gravity.  (See  pages  322  and  323.) 


SOLUTE    AND    SOLVENT.  33 

main  in  a  highly  supersaturated  condition  for  some  time,  especially  if 
the  solution  be  not  disturbed.  Separation  of  the  excess  can  usually  be 
induced  by  vigorous  shaking  or  by  adding  a  minute  particle  of  the 
solid  substance. 

If  the  separation  takes  place  comparatively  slowly,  large  regular- 
shaped  forms  result,  the  physical  properties  of  which  usually  vary 
with  the  direction  through  the  mass.  Such  forms  are  known  as  crystals 
and  the  process  of  separation  of  the  solid  from  the  solvent  as 
crystallization. 

Many  substances  in  separating  from  solution  in  crystal  form  take 
up  a  definite  amount  of  the  solvent  which  is  an  essential  constituent  of 
the  crystal  to  which  many  of  its  properties,  especially  its  color,  are 
due.  While  this  phenomenon  is  not  confined  to  any  particular  solvent, 
water  shows  especial  tendencies  in  this  direction  and  when  so  incor- 
porated in  the  crystalline  structure  is  known  as  water  of  crystallization. 

Water  of  crystallization  shows  none  of  the  ordinary  properties  and 
is  undoubtedly  combined  in  some  way  with  the  solid  substance.  The 
combination  is  less  stable  than  many  of  the  known  chemical  compounds, 
for  the  water  is  usually  driven  off  when  the  temperature  is  raised  much 
above  100°  and  then  shows  all  its  characteristic  properties. 

Under  ordinary  conditions  certain  crystals  tend  to  lose  their  water 
of  crystallization,  giving  it  up  to  the  atmosphere  surrounding  them.  Such 
substances  are  said  to  be  efflorescent.  On  the  other  hand  certain  sub- 
stances are  able  to  take  up  moisture  from  the  atmosphere  and  are  said 
to  be  hygroscopic.  The  absorption  and  condensation  of  water  from  the 
atmosphere  continues  in  some  cases  until  the  substance  passes  into 
solution,  this  phenomenon  being  known  as  deliquescence.  It  is  to  be 
noted,  however,  that  these  phenomena  depend  essentially  upon  the  amount 
of  water  vapor  present  in  the  atmosphere  and  the  same  substance  that 
is  efflorescent  in  a  dry  region  may  be  deliquescent  in  the  presence  of 
moisture. 

A  substance  which  is  free  from  water  is  said  to  be  anhydrous. 

If  the  separation  of  the  solute  from  the  solvent  takes  place  rapidly, 
the  product  is  finely  divided  and  is  called  a  precipitate,  the  process  being 
known  as  precipitation. 

The  precipitate  may  be  crystalline,  micro-crystalline  or  pulverulent 
(like  powder),  depending  on  the  size  of  the  crystals  formed  or,  if  the 
particles  are  so  small  as  to  be  apparently  without  definite  shape,  the 
precipitate  may  be  flocculent,  clotted  or  gelatinous.  In  every  case,  how- 
ever, the  solution  from  which  a  precipitate  separates  is  still  saturated 
with  the  precipitated  substance. 

4 


34  QUALITATIVE    ANALYSIS. 

If  it  be  desired  to  separate  a  substance  as  completely  as  possible  from 
solution,  conditions  must  be  brought  about  under  which  the  solubility 
of  the  substance  is  small ;  otherwise  much  of  the  solute  will  be  still  con- 
tained in  the  saturated  solution.  If,  for  instance,  it  is  desirable  to 
separate  barium  as  completely  as  possible  from  solution,  it  is  necessary 
to  decide  in  what  form  to  precipitate  it.  One  must  take  into  considera- 
tion the  facts  that  in  a  saturated  solution  of  barium  hydroxid  there  is 
1  part  of  the  solute  to  20  of  water,  of  barium  carbonate,  1  part  to  50,000 
of  water,  and  of  barium  sulphate,  1  to  400,000.  Hence  it  is  not  wise  to 
attempt  to  precipitate  as  the  hydroxid,  because  that  would  leave  behind 
one  part  of  hydroxid  in  every  20  parts  of  water,  but  rather  as  sulfate, 
which  will  leave  only  one  part  of  the  salt  in  400,000  of  water. 

Substances  which  have  extremely  small  solubilities  are  often  said  to 
be  insoluble  in  distinction  from  those  which  have  fairly  large  solubilities 
which  are  known  as  soluble.  There  are  no  substances  which  are  truly 
insoluble,  although  from  the  standpoint  of  qualitative  analysis  many 
are  practically  insoluble. 

The  term  precipitation  is  restricted  usually  to  substances  which  are 
fairly  insoluble,  while  the  word  crystallization  applies  to  those  which 
are  readily  soluble. 

Water  is  always  understood  to  be  the  solvent  to  which  reference 
is  made  when  no  other  solvent  is  specially  mentioned. 

Qualitative  analysis  consists  largely  of  a  study  of  comparative  solu- 
bilities. To  a  solution  which  may  contain  many  different  substances,  a 
reagent  is  added  often  capable  of  reacting  with  each.  Some  of  the 
reaction  products  are  so  slightly  soluble  that  the  first  trace  formed  will 
be  sufficient  to  saturate  the  liquid  and  the  remainder  will  then  separate 
as  an  insoluble  precipitate.  Other  reaction  products  will  have  such 
large  solubilities  that  the  total  amount  formed  will  not  be  sufficient  to 
saturate  the  solvent  present,  hence  all  will  remain  in  solution.  Because 
of  such  differences  in  solubility,  separations  are  readily  effected. 

Since  the  method  of  classifying  the  elements  into  groups,  according 
to  any  scheme  of  qualitative  analysis  depends  alone  on  the  varying 
solubilities  of  the  compounds  formed,  it  does  not  necessarily  mean  that 
those  elements  that  most  closely  resemble  each  other  in  their  chemical 
properties  and  their  compounds  will  be  found  in  the  same  group.  The 
metals  mercury,  cadmium,  zinc  and  magnesium  closely  resemble  each 
other,  yet  are  found  two  of  them  in  Group  I,  one  in  Group  III  and  one 
in  Group  V,  owing  chiefly  to  very  varying  solubilities  of  their  respective 
sulfids.  The  grouping  of  elements  in  qualitative  analysis  denotes 


SOLUTE    AND    SOLVENT.  35 

similar  solubilities  of  similar  salts  of  the  different  metals  in  any  group 
and  nothing  further. 

The  most  satisfactory  classification  of  the  elements  into  groups  based 
on  chemical  similarity  in  properties  is  known  as  the  Periodic  System  and 
is  given  on  pages  64  and  65. 

An  interesting  case  of  the  interworking  of  the  laws  governing  solu- 
tion and  ionization  may  be  developed  from  a  consideration  of  the 
phenomena  resulting  when  substances  insoluble  in  water  are  treated  with 
solutions  of  acids.  The  insoluble  salts  of  weak  acids  dissolve  in  solu- 
tions of  strong  acids,  but  the  insoluble  salts  of  strong  acids  will  not 
dissolve  in  the  solutions  of  other  strong  acids. 

When  calcium  phosphate  comes  in  contact  with  water,  it  forms  a  sat- 
urated solution.  Practically  all  of  the  dissolved  salt  exists  in  the 
ionic  condition,  for  the  dilution  is  large  since  the  solubility  is  small. 
There  is  equilibrium  between  the  ions  and  the  undissociated  salt  and 
between  the  dissolved  and  the  undissolved  portions  as  well.  If  any 
strong  acid  be  added,  it  introduces  a  relatively  high  concentration  of 
hydrogen  ions.  This  factor  taken  into  consideration  with  the  concentra- 
tion of  phosphate  ions  introduced  by  the  dissolved  salt  is  too  great  for 
equilibrium  between  dissociated  and  undissociated  phosphoric  acid  since 
this  acid  dissociates  to  a  small  degree  only.  By  the  equilibrium  reaction 
undissociated  phosphoric  acid  is  formed,  the  concentration  of  phosphate 
ions  is  lessened  and  the  equilibrium  between  the  dissolved  and  undis- 
solved portion  of  the  salt  is  disturbed.  More  salt  passes  into  solution 
to  restore  the  original  concentration  of  phosphate  ions.  This  causes  the 
formation  of  more  undissociated  phosphoric  acid  and  a  second  adjust- 
ment of  concentrations  becomes  necessary.  In  this  way  the  slightly 
soluble  salt  dissolves  readily  as  long  as  there  is  a  high  concentration  of 
hydrogen  ions  in  the  solution. 

There  remains  in  solution  an  abundance  of  calcium  kathions  and 
the  anions  of  the  acid  added.  The  phosphoric  acid  is  for  the  most  part 
undissociated.  By  the  application  of  the  principles  of  equilibrium  and 
mass  action  alone  it  is  thus  possible  to  explain  how  the  "stronger  acid 
takes  the  base  from  the  weaker  acid. ' ' 

Insoluble  hydroxids  are  all  readily  soluble  in  acids  and  are  to  be 
considered  as  salts  of  the  very  weak  acid,  HOH.  Many,  probably  all, 
oxids  unite  with  water  to  form  hydroxids  to  some  extent  at  least.  Their 
solubility  in  dilute  acid  solutions  is  to  be  explained  exactly  as  the  solu- 
bility of  the  hydroxids. 

On  the  other  hand,  solution  of  the  very-slightly-soluble  silver  chlorid 


36  QUALITATIVE    ANALYSIS. 

is  not  appreciably  aided  by  the  addition  of  the  strong  nitric  acid, 
although  both  of  the  products  of  the  possible  metathesis,  silver  nitrate 
and  hydrochloric  acid,  are  extremely  soluble.  The  conditions  of  equi- 
librium between  the  ionized  and  non-ionized  portions,  and  the  dissolved 
and  undissolved  portions  of  the  silver  chlorid  are  not  disturbed  by  the 
highest  concentrations  of  hydrogen  ions  obtainable.  There  is  no  tendency 
toward  reduction  in  the  concentration  of  any  of  the  ions  present  in  solu- 
tion by  the  formation  of  undissociated  substances  since  all  the  compounds 
possible  of  formation  dissociate  to  a  very  high  degree.  Hence  the  system 
maintains  the  equilibrium  which  it  first  reached. 

The  hydroxids  of  lead,  tin,  zinc  and  aluminum,  substances  insoluble 
in  water,  furnish  another  application  of  the  same  principle  by  their 
solubility  in  strong  bases.  Here  water  plays  the  part  of  a  weak 
acid  (see  pages  29  and  176).  In  the  presence  of  a  high  concentration 
of  hydroxyl  ions  from  sodium  or  potassium  hydroxid  (ammonium  hy- 
droxid  is  not  sufficiently  dissociated),  lead  hydroxid,  Pb02H2,  acts 
as  an  acid  and  separates  hydrogen  ions  (see  pages  22  and  143).  Water 
is  formed  from  these  components  and  the  initial  equilibrium  dis- 
turbed. The  lead  hydroxid  continues  to  dissolve  and  water  continues 
to  form  until  all  the  lead  hydroxid  has  passed  into  solution.  There 
is  left  an  abundance  of  the  ions  Pb02  and  Na  to  form  the  salt,  sodium 
plumbite,  Na2Pb02. 

In  general,  whenever  any  reagent  is  added  to  a  saturated  solution  of 
an  insoluble  substance,  the  solution  being  in  contact  with  an  excess  of  the 
solute,  if  the  added  reagent  causes  the  concentration  of  the  ions  of  the 
solute  to  become  less,  the  insoluble  substance  tends  to  dissolve. 

SEPARATION   OF    SOLIDS   FROM   LIQUIDS. 

A  liquid  may  be  separated  from  a  substance  dissolved  in  it  by 
causing  the  liquid  to  become  volatile  or  to  pass  into  the  gaseous  state. 
Such  a  method  of  separation  is  not  applicable  when  both  solute  and  sol- 
vent are  volatile  under  the  same  conditions.  If  the  liquid  is  not  retained, 
the  process  is  known  as  evaporation.  At  high  temperatures  evapora- 
tion takes  place  more  rapidly  than  at  low  temperatures.  At  ordinary 
temperatures  it  is  said  to  be  spontaneous. 

If,  in  the  separation  of  solute  from  solvent,  the  liquid  be  collected, 
the  process  is  called  distillation.  The  liquid  is  usually  caused  to  evapo- 
rate rapidly  by  applying  heat.  The  vapors  are  liquefied  again  by 
passing  through  a  cooling  apparatus  called  a  condenser.  The  con- 
densed liquid,  thus  separated  from  all  non-volatile  substances,  is  known 
as  the  distillate. 


SEPARATION    OF    SOLIDS   FROM   LIQUIDS.  37 

A  liquid  may  be  separated  from  substances  suspended  in  it  by 
straining  it  through  a  porous  membrane,  the  size  of  the  pores  being  such 
that  while  the  particles  of  the  liquid  are  readily  allowed  to  pass  through, 
the  passage  of  the  suspended  matter  is  prohibited.  This  membrane  is 
known  as  the  filter;  the  liquid  passing  through  its  pores  is  called  the 
filtrate  and  the  suspended  matter  retained  by  the  filter  is  termed  the 
residue  or  precipitate. 

The  process  of  filtration  is  merely  a  mechanical  screening  out  of 
the  suspended  matter  because  the  size  of  these  particles  is  very  much 
larger  than  the  molecules  of  the  liquid.  Obviously  the  size  of  the 
particles  to  be  removed  will  determine  how  small  the  diameter  of  the 
pores  in  the  filter  must  be  to  effect  a  separation.  Almost  any  powder 
which  can  be  obtained  mechanically  may  be  satisfactorily  separated  by 
ordinary  filter  paper. 

Just  as  the  particles  of  a  solid  substance  move  over  each  other 
with  friction,  so  do  the  molecules  of  a  liquid  move  with  friction  on  each 
other.  Friction  between  molecules  as  distinguished  from  friction  between 
masses  is  called  internal  friction,  or,  more  often  in  the  case  of  liquids, 
viscosity.  Since  heat  increases  the  freedom  of  motion  of  the  molecules 
of  a  body,  the  viscosity  of  a  liquid  is  lessened.  For  this  reason  it  will 
pass  through  a  filter  much  more  rapidly  when  hot  than  when  cold.  It  is 
expedient,  therefore,  to  filter  all  liquids  as  hot  as  they  can  be  con- 
veniently handled  unless  there  is  some  definite  reason  for  doing  otherwise. 

By  certain  chemical  reactions  it  is  possible  to  produce  precipitates 
so  finely  divided  that  they  will  pass  through  the  finest  filter.  The  par- 
ticles are  so  minute  as  to  be  invisible  to  the  most  powerful  micro- 
scope. The  liquid  in  which  they  are  suspended,  appears  perfectly  clear 
and  has  all  the  appearance  of  a  solution,  although  its  properties  are 
those  of  the  pure  solvent.  Such  a  finely  divided  precipitate  is  known 
as  a  colloidal  precipitate  or  as  a  pseudo-solution. 

On  boiling  a  pseudo-solution  the  particles  generally  coalesce,  the 
liquid  becomes  turbid  and  may  then  be  filtered  in  the  ordinary  way.  By 
the  addition  of  a  strongly  dissociating  substance  to  the  pseudo-solution 
a  visible  precipitation  may  often  be  effected. 

Other  precipitates  are  frequently  met,  intermediate  in  size  between 
the  colloidal  precipitate  and  that  which  may  be  readily  filtered.  Some- 
times these  may  be  so  finely  divided  as  to  run  through  the  filter  paper. 
This  may  often  be  prevented  by  using  two  filters  instead  of  one  or  by 
pouring  the  cloudy  filtrate  back  onto  the  paper  repeatedly.  The  pores 
in  the  filter  will  gradually  become  filled  with  the  precipitate  to  such 


38  QUALITATIVE    ANALYSIS. 

an  extent  that  the  filtrate  will  finally  run  through  clear.  By  allowing 
a  finely  divided  precipitate  to  stand  for  some  time  in  hot  water,  the 
particles  often  coalesce  and  become  more  crystalline  after  which  they 
may  be  readily  filtered. 

Sometimes  a  precipitate,  on  standing  in  contact  with  a  cold  solution, 
will  exhibit  a  tendency  toward  the  formation  of  colloidal  precipitates. 
The  result  is  the  formation  of  a  slime  which  clogs  the  filter  and  makes 
the  process  of  filtration  exceedingly  slow  and  tedious.  As  the  tendency 
toward  the  formation  of  colloidal  solutions  decreases  as  the  tempera- 
ture rises,  rapid  filtration  is  assured  by  the  use  of  hot  liquids. 

The  regulation  round  filter  should  be  folded  twice  to  form  a  quad- 
rant and  then  opened  so  as  to  form  a  cone  with  three  thicknesses  of 
paper  on  one  side  and  one  on  the  other.  In  this  shape  it  should  be  in- 
serted in  the  funnel  and  the  paper  fitted  to  the  glass  by  adjusting  the 
folds  a  trifle  if  necessary.  Moisten  with  water  and  press  firmly  against 
the  sides  of  the  funnel  so  that  the  filter  and  funnel  are  everywhere  in 
contact.  When  a  solution  is  poured  upon  a  properly  fitted  filter,  the 
stem  of  the  funnel  will  fill  with  the  hot  filtrate,  the  weight  of  which 
will  tend  to  suck  the  liquid  through  the  paper  and  make  the  process 
more  rapid  than  if  the  filter  is  laid  loosely  in  the  funnel. 

In  order  to  separate  the  precipitate  completely  from  the  filtrate,  it 
should  be  allowed  to  drain  as  dry  as  possible,  after  which  it  is  necessary 
to  wash  it.  Hot  water  should  ordinarily  be  used,  since  a  small  volume 
will  dissolve  any  adhering  substances  to  a  greater  degree  than  the 
same  volume  of  cold  water.  For  the  reasons  already  considered,  washing 
with  hot  water  tends  to  keep  the  precipitate  in  better  condition  than 
cold  water.  After  each  washing,  the  precipitate  should  be  allowed  to 
drain  thoroughly  before  more  water  is  poured  onto  the  paper. 

If  the  precipitate  settles  rapidly  and  it  is  undesirable  for  any 
reason  to  separate  it  by  filtration,  the  liquid  may  be  poured  off  as 
completely  as  possible  and  hot  water  added  and  the  mixture  shaken. 
After  allowing  the  precipitate  to  settle  again  the  wash  water  may  be 
poured  off  and  the  precipitate  treated  as  before.  This  process  is  known 
as  washing  by  decantation. 

OXIDATION   AND    REDUCTION. 

When  heated  in  air  or  an  atmosphere  of  oxygen,  many  substances 
gain  in  weight.  Because  of  the  chemical  activity  of  oxygen  at  high 
temperatures,  they  unite  with  it  and  form  oxids.  This  process  is  known 
as  oxidation. 


OXIDATION    AND    SEDUCTION.  39 

On  the  other  hand,  when  oxids  are  heated  in  an  atmosphere  of  hy- 
drogen, they  lose  their  oxygen  since  it  unites  with  hydrogen  to  form 
water.  This  process  is  called  reduction  and  is  the  reverse  of  oxidation. 

Certain  non-metallic  elements,  other  than  oxygen  but  equally  reactive, 
behave  in  a  very  analogous  manner  under  the  same  conditions.  Hence 
it  is  customary  to  extend  these  conceptions  and  to  consider  as  oxidation 
any  process  in  which  a  non-metallic  element  is  added  to  a  substance  or 
hydrogen  is  removed  from  it.  Reduction  is  any  process  in  which  a  non- 
metallic  element  is  taken  away  from  a  substance  or  hydrogen  is  added 
to  it. 

That  the  removal  of  hydrogen  is  equivalent  to  the  addition  of  oxy- 
gen, and  vice  versa,  can  best  be  shown  by  a  concrete  illustration.  If 
ammonia,  NH3,  be  oxidized,  we  may  consider  that  the  oxygen  attaches 
itself  directly  forming  a  compound,  NH3O.  If  this  new  compound  sub- 
sequently decomposes  into  NH  and  H20,  it  does  not  lessen  the  fact  that 
NH  is  an  oxidation  product  of  NH3  as  well  as  H20.  Yet  NH  is  simply 
NH3  from  which  hydrogen  has  been  removed.  Further  oxidation  may 
remove  all  the  hydrogen  and  later  add  oxygen.  In  the  oxidation  of 
ammonia  all  of  the  following  steps  are  conceivable  and  all  are  known 
with  the  exception  of  the  third : 

2NH3    N2H4    N2H2    N2    N2O    N202    N2O3    N2O4    N2O5 

Looked  at  from  the  other  direction,  each  one  of  these  bodies  may 
be  regarded  as  a  reduction  product  of  N2O5,  in  which  case  reduction  is 
seen  to  be  no  less  a  removal  of  oxygen  than  the  addition  of  hydrogen. 
N205  is  to  be  regarded  as  the  highest  oxidized  product  of  ammonia  and 
NH3  as  the  lowest  reduction  product  of  nitrogen  pentoxid. 

Any  substance  which  causes  oxidation  is  called  an  oxidizing  agent. 
Any  substance  causing  reduction  is  known  as  a  reducing  agent. 

Oxidation  and  reduction  are  reciprocal  terms.  One  process  can  not 
take  place  without  the  other.  If  one  substance  gives  up  its  non-metallic 
constituents  wholly  or  in  part  to  another  substance,  the  latter  is  oxidized, 
but  the  oxidizing  agent  is  itself  reduced.  Hence  any  substance  which  is 
readily  oxidized  is  a  reducing  agent  and  any  substance  which  is  readily 
reduced  is  an  oxidizing  agent. 

Under  these  circumstances  it  is  not  surprising  to  find  that  certain 
substances  may  be  either  oxidizing  or  reducing  agents  depending  on 
the  conditions  present.  In  the  presence  of  a  strong  oxidizing  agent  they 
may  act  as  reducing  agents,  while  in  the  presence  of  a  strong  reducing 
agent  they  may  act  as  oxidizing  agents. 


40  QUALITATIVE    ANALYSIS. 

Those  elements  which  are  ordinarily  used  as  oxidizing  agents  are 
oxygen,  chlorin,  bromin  and  iodin,  together  with  compounds  from  which 
these  elements  may  be  readily  liberated. 

Hydrogen,  carbon  and  Compounds  containing  these  elements  in  con- 
siderable proportions  are  the  ordinary  reducing  agents.  Sulfur  and 
its  compounds  are  used  as  reducing  agents  also,  but  the  element  and 
its  dioxid  may  act  as  oxidizing  agents  in  the  presence  of  stronger  re- 
ducing agents. 

By  far  the  most  active  reducing  agent  known  is  hydrogen,  especially 
at  the  moment  when  it  is  liberated  from  a  compound.  If  a  metal  be 
thrown  into  a  solution  of  an  acid,  any  substance  present  in  solution  with 
the  acid  will  be  reduced  more  quickly  than  if  the  hydrogen  is  collected 
and  then  bubbled  through  a  similar  solution.  When  in  this  very  active 
state,  it  is  said  to  be  in  the  nascent  condition. 

Possibly  the  most  satisfactory  explanation  that  has  been  offered  of 
this  phenomenon  is  concerned  with  the  structure  of  the  molecule.  The 
molecule  of  gaseous  hydrogen  is  known  to  consist  of  at  least  two  atoms 
joined  to  each  other.  This  structure  may  be  represented  graphically  by 
the  formula  H  —  H,  each  atom  being  linked  to  its  fellow  by  a  single 
bond  since  hydrogen  is  a  univalent  element.  For  an  instant  after  its 
liberation  from  a  compound  it  is  thought  to  exist  in  the  atomic  condition, 
the  atoms  being  not  yet  joined  together.  This  condition  may  be  repre- 
sented by  the  symbol  H  — ,  the  free  bond  indicating  that  the  atom  exists 
alone  uncombined  with  any  other.  It  is  easily  conceivable  that  any 
substance  would  be  more  active  when  its  atoms  were  free  than  when 
joined  together  in  pairs  or  otherwise  to  form  molecules,  since  in  the 
latter  condition  the  molecule  must  be  broken  up  before  its  atoms  can 
unite  with  the  atoms  of  other  substances. 

Elements  other  than  hydrogen  exhibit  similar  phenomena.  The 
nascent  condition  of  an  element  may  be  indicated  by  free  bonds  attached 
to  the  symbol  for  its  atom,  one  bond  being  added  for  each  valence  shown. 
Nascent  chlorin,  being  univalent,  is  represented  by  the  symbol  Cl  — 
nascent  oxygen,  being  divalent,  is  represented  by  the  symbol  0=. 

Hydrogen  is  the  best  reducing  substance  known,  yet  a  compound  in 
which  hydrogen  is  present  to  only  a  small  proportion  may  be  a  more  active 
reducing  agent  than  hydrogen  gas,  provided  it  can  be  made  to  liberate 
its  hydrogen  in  the  nascent  condition.  Similarly  oxygen  gas  is  not  nec- 
essarily the  best  oxidizing  agent,  for  any  substance  which  will  liberate 
nascent  oxygen  will  be  a  more  active  oxidizing  agent.  Hence  it  is  that 


OXIDATION    AND    KEDUCTION.  41 

many  compounds  are  known  which  are  more  active  oxidizing  agents 
than  air  or  even  pure  oxygen. 

In  the  reactions  with  which  qualitative  analysis  is  concerned,  the 
following  substances  act  as  oxidizing  agents: 

(a)   Oxygen      (6)   Nitrous  acid,  HNO2  (c)   Stannic  salts — from 

Sulfur  Nitric  acid,  HNO3  the  oxid  Sn02 

Chlorin  Hypochlorous  acid,  HC10  Ferric    salts — from 

Bromin  Chloric  acid,  HC103  the  oxid  Fe=0 

lodin  Permanganic  acid,  HMn04  \ 

Chromic  acid,  H2Cr04*  ° 

Dichromic  acid,  H2Cr207  -p  _Q 

Hydrogen  peroxid,  H202  Mercuric    salts- 

from     the     oxid 
HgO 

(a)  The  elementary  substances  usually  show  their  oxidizing  action 
by  direct  addition  to  the  reducing  substances,  as  the  following  reactions 
indicate : 

(1)  S02  +  Ot-+S03 

(2)  SnS  +  S- 

(3)  FeCl2-f 

(4)  SnCl2  +  Br2  -^  SnCl2Br2  or,  doubling  the  quantities, 

" 

The  halogen  elements,  cnlorm,  broTmrT  and  iodin,  frequently  act  as 
oxidizing  agents  when  it  is  preferable  to  indicate  their  action  as  an 
addition  of  oxygen  rather  than  of  halogen.  In  such  cases  it  is  customary 
to  consider  that  the  halogens  act  on  water,  decomposing  it,  forming  a 
hydrogen  halogen  acid  and  liberating  nascent  oxygen  which  then  adds 
directly  to  the  reducing  agent.  The  following  reactions  will  clearly 
indicate  the  nature  of  the  process. 

2CI  +  H2O  ->  2HC1  +  0  =         HN02  +  0  =  ->  HNO3 

*  Chromic  acid  is  practically  unknown,  since  it  loses  water  and  passes  into 
dichromic  acid  according  to  the  reaction : 

2H2CrO<  -»  H2Cr207  +  H2O 

It  is  included  in  the  table  because  its  salts  are  well  known. 

t  Although  it  is  a  well-established  fact  that  the  molecules  of  non -metallic  elements 
consist  of  at  least  two  atoms,  no  attempt  will  be  made  to  introduce  at  least  a  molecule 
of  all  elements  into  equations  when  simplicity  is  to  be  gained  by  representing  only 
a  single  atom. 


42  QUALITATIVE   ANALYSIS. 

Both  reactions  may  be  combined  in  one. 

HN02  +  2CI  +  H20  ->  HN03  +  2HC1 

(&)  All  of  this  group  of  oxidizing  agents  are  acids.  Salts  of  these 
acids,  especially  the  sodium  and  potassium  salts,  are  frequently  used 
instead  of  the  acids.  The  lead  and  manganese  salts  of  hydrogen 
peroxid,  lead  and  manganese  dioxids,  are  quite  as  commonly  used 
as  sodium  peroxid.  Since  a  salt  is  generally  more  stable  than  the 
corresponding  acid,  as  an  oxidizing  agent  a  salt  is  less  energetic  than 
the  free  acid. 

For  the  purpose  of  simplifying  oxidation  and  reduction  reactions  it 
is  often  advisable  to  consider  an  acid  or  its  salt  as  composed  of  an  oxid 
of  a  metallic  element  and  an  oxid  of  a  non-metallic  element,  hydrogen 
being  considered  as  a  metallic  element.  Chromic  acid,  H2Cr04,  is 
equivalent  to  H20  and  Cr03;  potassium  chromate  to  K2O  and  Cr03. 
None  of  the  oxygen  available  in  a  reaction  comes  from  water  or  the 
metallic  oxid  corresponding  to  it,  as  K2O  or  Na20,  hence  no  error  is 
made  in  disregarding  these  factors.  In  this  way  we  may  consider 

Nitrous  acid,  2HN02  as  H20  +  N203 

Nitric  acid,  2HN03,  as  H20  +  N2O5 

Hypochlorous  acid,  2HC10,  as  H20  +  C12O 

Chloric  acid,  2HC103,  as  H2O  +  C12O5 

Permanganic  acid,  2HMn04,  as  H2O  +  Mn2O7 

Chromic  acid,  H2Cr04,  as  H20  +  Cr03 

Dichromic  acid,  H2Cr2O7,  as  H2O  +  2Cr03 

Hydrogen  peroxid,  H202,  as  H20  +  O 

Since  these  non-metallic  oxids  are  acids  minus  water,  they  are  known 
as  acid  anhydrids. 

That  this  series  of  acids  and  salts  act  as  oxidizing  agents  is  due 
to  the  fact  that  ordinarily  in  the  presence  of  reducing  agents  their 
anhydrids  are  not  as  stable  as  some  lower  form  containing  less  oxygen. 
Thus  N2O2  is  the  oxid  of  nitrogen  usually  formed  as  a  reduction  product, 
hence  both  nitrous  and  nitric  acids  break  down  into  N202,  water  and 
nascent  oxygen.  The  amount  of  oxygen  liberated  depends  on  the  differ- 
ence in  the  oxygen  content  between  the  oxidizing  agent  and  the  com- 
pound stable  in  the  presence  of  reducing  agents.  By  the  application  of 
this  principle  an  oxygen  equivalent  may  be  calculated  for  each  oxidizing 
agent.  The  conditions  of  the  experiment  influence  the  amount  of  oxygen 
liberated  somewhat,  but  the  equivalents  given  by  the  following  equa- 
tions may  be  considered  as  typical. 


OXIDATION    AND    REDUCTION.  43 

Nitrous  acid  (2  mols.),  N2O3  -^    N2O2  +     O  = 

Nitric  acid*   (2  mols.),  N2O5  ->    N202  +30  = 

Chromic  acid  (2  mols.),  2Cr08          ->    Cr203  +30  = 

Dichromic  acid  (1  mol.),  2Cr03         -»    Cr203  +30  = 

Permanganic  acid  (2  mols.),  Mn207-*  2Mn02  +  30=  in  neutral  or  al- 
kaline solution 

Permanganic  acid  (2  mols.),  Mn2O7->   2MnO  +  5O  =  in  acid  solution 
Hypochlorous  acid  (2  mols.),  C12O    -»  0=  +  2Clf 

Chloric  acid  (2  mols.),  C1205  ->  50  =  +  2Clf 

Hydrogen  peroxid  (1  mol.),  0          —>  O  = 

By  the  intervention  of  a  molecule  of  water  in  those  reactions  in 
which  the  direct  addition  of  chlorin,  bromin  or  iodin  is  not  to  be  indi- 
cated, the  action  of  the  halogen  elements  may  be  referred  to  an  oxygen 
basis  and  their  oxygen  equivalent  indicated  as  follows: 

C12  +  H20  ->  2HC1  +  0  = 
Br2  +  H~0  ->  2HBr  +  0  = 
I2  +  H20-»2HI   +0=1 

*  One  of  the  commonest  oxidizing  agents  is  known  as  aqua  regia.  It  is  made 
by  mixing  nitric  and  hydrochloric  acids  and  warming.  Since  hydrochloric  acid  is  not 
an  oxidizing  agent,  the  oxidizing  power  of  aqua  regia  must  be  due  entirely  to  nitric 
acid.  In  the  proportions  of  three  of  hydrochloric  acid  to  one  of  nitric  acid  the  fol- 
lowing reaction  takes  place: 

3HC1  +  HNO3  -»  3C1  —      +O=     +N  +  2H2O 

Since  nascent  chlorin  attacks  many  mineral  substances  far  more  readily  than 
nascent  oxygen,  aqua  regia  is  a  more  general  oxidizing  agent  than  nitric  acid  alone. 
In  this  combination  nitric  acid  is  broken  down  to  free  nitrogen,  thus  a  greater  equiva- 
lent of  oxygen  is  obtained  than  is  usual  with  nitric  acid. 

Aqua  regia  consists  of  a  mixture  of  hydrochloric  acid,  nitric  acid,  nascent 
chlorin  and  nascent  oxygen.  Its  action  may  be  more  satisfactorily  expressed  by  these 
individual  reagents,  such  as  are  needed,  than  by  attempting  to  write  an  equation  that 
should  include  the  whole  mixture.  The  oxidation  of  mercurous  chlorid  by  aqua  regia 
is  best  expressed  by  the  equation : 

(HgCl)2  +  2CZ-     -»     2HgCl2 
The  oxidation  of  nickel  sulfid  uses  other  reagents: 

NiS  +  40  =     +  2HC1  ->  NiCL  +  H2SO4 

t  Although  hypochlorous  and  chloric  acids  decompose  in  this  way,  their  salts  do 
not.  The  latter  break  down  into  chlorids  and  nascent  oxygen,  thus: 

KC1O  -»  KC1  +  0  =      and       KC1O3  _»  KC1  +  30  = 

The  oxygen  equivalents  of  salt  and  acid  are  the  same,  the  only  difference  being  that 
the  salts  liberate  oxygen  only  while  the  acids  liberate  both  oxygen  and  chlorin. 

$  This  reaction  with  iodin  is  not  known  to  take  place  to  an  appreciable  extent 
under  ordinary  conditions,  but  for  the  sake  of  systematization  it  may  be  allowed  to 
appear  in  this  connection. 


44  QUALITATIVE    ANALYSIS. 

(c)  Since  the  salts  of  any  metal  may  be  formed  by  the  action  of 
acids  on  the  various  oxids  of  the  metal,  these  oxids  offer  a  simple  basis 
for  the  consideration  of  oxidation  and  reduction  reactions.  Any  bi- 
valent salt  of  iron  may  be  formed,  theoretically  at  least,  by  dissolving 
the  bivalent  ferrous  oxid,  FeO,  in  the  necessary  acid.  Similarly,  any 
trivalent  salt  of  iron  may  be  formed  by  the  action  of  the  necessary  acid 
on  the  trivalent  ferric  oxid,  Fe203.  The  oxidation  of  ferrous  chlorid, 
FeCL,  ferrous  sulfate,  FeS04,  or  ferrous  acetate,  Fe(C2H302)2  to  ferric 
chlorid,  FeCl3,  ferric  sulfate,  Fe2(S04)3,  or  ferric  acetate,  Fe(C2H302)3, 
may  be  considered  as  the  oxidation  of  ferrous  oxid  to  ferric  oxid  and 
the  action  of  the  necessary  acids  on  the  oxidation  product.  By  this  means 
the  oxidation  of  any  particular  salt  is  referred  to  a  general  basis  and 
the  equations  much  simplified.  For  instance,  let  ferrous  sulfate  be 
oxidized  by  the  oxygen  of  the  air.  To  oxidize  ferrous  oxid  to  ferric 
oxid  requires  one  atom  of  oxygen  for  every  two  molecules  of  ferrous 
oxid  treated,  for 

Fe=0 

0         0->         0 


One  atom  of  oxygen  will  oxidize  two  molecules  of  ferrous  sulfate  in 
the  same  way  : 

Fe=S04 

\ 

2Fe  =  S04  +  0-^          0 

Fe=S04 

This  compound  may  be  considered  as  a  mixture  of  ferric  sulfate  and 
ferric  oxid  in  the  proportion  of  2Fe2(S04)3  to  !Fe203.  Evidently  there 
is  need  of  additional  sulfuric  acid  to  convert  it  entirely  into,  ferric 

sulfate,  for 

Fe=S04      H  Fe=S04 

\  \  \ 

0-j-         S04^          S04  +  H2O 

/  /  / 

Fe=S04      H  Fe=SO4 

Ferrous  sulfate  can  not  be  oxidized  entirely  over  into  ferric  sulfate 
except  in  the  presence  of  free  sulfuric  acid. 

From  a  consideration  of  this  and  other  similar  reactions  it  is  seen 
that  acid  is  used  up  in  many  oxidation  reactions.     For  this  reason 


OXIDATION    AND    EEDUCTION.  45 

it  is  generally  true  that  oxidation  takes  place  better  in  acid  than  in 
neutral  or  alkaline  solutions. 

From  the  reactions  already  considered  it  will  be  apparent  that  the 
element  oxidized  generally  has  its  valence  increased.  Divalent  iron 
becomes  trivalent,  divalent  tin  becomes  tetravalent.  Divalent  sulfur  in 
hydrogen  sulfid,  H2S,  oxidizes  to  the  hexavalent  condition  in  sulfuric 
acid,  H2SO4.  For  each  atom  of  oxygen  or  its  equivalent  added  to  an 
element,  its  valence  is  increased  by  two.  Nevertheless  increase  of  valence 
is  not  always  brought  about  by  oxidation.  Hydrogen  gas  burns  to  water, 
marsh  gas  to  water  and  carbon  dioxid.  The  following  equations  repre- 
sent the  reactions  taking  place  and  show  the  valences  of  the  elements  in 
each  compound: 

H  —  H     +     0    ->     H  — 0  — H 

H  O 

H  —  C  —  H    +     202    ->    2H  — 0  — H     +     C 

I  \ 

H  0 

Some  elements  form  a  hydroxid  corresponding  to  each  of  the  different 
valences  shown  by  the  element.  The  valence  of  manganese  changes 
from  two  to  seven  and  most  of  the  corresponding  hydroxids  are  known. 
The  hydroxid  of  divalent  manganese  acts  as  a  fairly  strong  base  and 
forms  salts,  such  as  manganous  chlorid,  MnCl2,  in  which  manganese 
plays  the  part  of  a  metal.  The  hydroxid  of  manganese  showing  seven 
valences,  on  the  other  hand,  acts  as  a  strong  acid  and  forms  salts,  such 
as  potassium  permanganate,  KMn04,  in  which  manganese  appears  in 
the  anion.  The  hydroxids  of  manganese  showing  a  valence  of  three, 
four  and  six  are  intermediate  in  properties,  trivalent  manganese  acting 
as  a  weak  base,  tetravalent  manganese  as  a  very  weak  acid.  Hexavalent 
manganese  forms  an  acid  which  is  stronger  than  the  corresponding 
compound  of  tetravalent  manganese,  but  not  so  strong  as  permanganic 
acid.  (See  pages  212-220.) 

While  the  hydroxids  of  sulfur  are  all  of  an  acid  nature,  sulfurous 
acid  corresponding  to  the  tetravalent  condition  is  much  weaker  than 
the  hexavalent  sulfur  in  sulfuric  acid.  Persulfuric  acid  in  which  sulfur 
shows  a  valence  of  seven  is  still  stronger.  (See  pages  184-190.) 

Thus  it  is  seen  that  when  an  element  is  oxidized,  with  an  increase  in 
valence  there  is  a  decrease  in  the  basic  properties  and  a  corresponding 
increase  in  the  acid  properties  of  the  element  as  is  shown  by  the  nature 
of  the  compounds  formed.  For  this  reason  the  elements  forming  the 


46  QUALITATIVE    ANALYSIS. 

strongest  bases  are  all  univalent;  divalent  bases  are  less  strong.  Bases 
from  tri-  and  tetravalent  elements  are  decidedly  weaker.  So  weak  are 
they  that  their  salts  generally  hydrolyze  in  solution.  For  these  reasons 
tri-  and  tetravalent  kathions  are  not  so  well  known  as  those  of  a  smaller 
valence.  (See  pages  11  and  20.)  Conversely,  a  comparison  shows  that 
if  an  element  is  capable  of  forming  more  than  one  oxygen*  acid,  the 
most  highly  oxidized  compound  is  the  strongest  acid.\ 

On  the  other  hand,  reduction  tends  to  reduce  acid  properties  and  to 
develop  basic  properties.  The  hydroxid  in  which  chromium  shows  seven 
valences  is  the  very  strong  perchromic  acid.  Hexavalent  chromium  forms 
the  strong  chromic  acid.  Trivalent  chromium  hydroxid  acts  as  a  very 
weak  base.  (See  pages  191-195.) 

The  very  great  difference  in  the  properties  of  the  compounds  formed 
by  an  element  in  a  highly  oxidized  condition  and  the  properties  of  the 
compounds  formed  by  the  same  element  in  a  reduced  condition  must  be 
constantly  borne  in  mind  in  making  an  analysis.  After  treatment  with 
strong  oxidizing  agents,  manganese  and  chromium  must  be  sought  for 
among  the  acids.  They  will  be  found  among  the  bases  after  the  intro- 
duction of  reducing  agents.  (See  pages  263  and  265.) 

In  many  cases  a  reducing  agent  acts  more  efficiently  in  an  alkaline 
solution  than  in  the  presence  of  an  acid.  For,  as  oxidation  tends  to  de- 
crease the  amount  of  free  acid  present  by  increasing  the'  valence  of 
metallic  elements,  thus  giving  them  a  greater  power  of  combining  with 
acids  (see  page  44),  so,  conversely,  reduction  tends  to  develop  free  acid 
by  decreasing  the  valence  of  metallic  elements,  lessening  their  power  of 
combining  with  acid  radicles  and  therefore  setting  acid  free. 

The  common  substances  which  act  as  reducing  agents  in  qualitative 
analysis  are: 

Hydrogen  Hydrocyanic  acid,  HCN    • 

Carbon  Thiocyanic  acid,  HSCN 

Sulfur  Nitrous  acid,  HN02 

Alcohol,  H6C2O  Hydrochloric  acid,  HC1 

Oxalic  acid,  H2C2O4  Hydrobromic  acid,  HBr 

Tartaric  acid,  H6C4O6  Hydriodic  acid,  HI 

Hydrosulfuric  acid,  H2S      Stannous  salts  —  from  the  oxid  Sn  =  0 

Sulfurous  acid,  H2SO3          Ferrous  salts  —  from  the  oxid  Fe  —  0 

Thiosulfuric  acid,  H2S203 

*  Acids  which  do  not  contain  oxygen  are  not  included  in  this  generalization. 

f  The  element  phosphorus  offers  an  exception  to  the  rule,  phosphorous  acid,  a 
compound  of  trivalent  phosphorus,  being  stronger  than  the  pentavalent  compound, 
phosphoric  acid. 


OXIDATION    AND    EEDUCTION.  47 

That  these  substances  act  as  reducing  agents  is  due  to  the  fact  that 
in  the  presence  of  oxidizing  agents  these  compounds  are  not  as  stable 
as  some  more  highly-oxidized  forms.  Consequently  they  take  up  as 
much  oxygen  as  is  necessary  to  convert  them  into  compounds  stable  in 
the  presence  of  oxidizing  agents.  One  substance  may  not  absorb  the  same 
•amount  of  oxygen  as  another.  For  this  reason  the  values  of  reducing 
agents  vary  considerably.  The  oxidation  products  depend  somewhat 
upon  the  conditions  of  the  experiment,  consequently  the  efficiency  of  a 
reducing  agent  may  vary  under  different  conditions.  The  reactions 
given  below  indicate  the  amount  of  oxygen  taken  up  by  each  substance 
under  conditions  which  may  be  considered  typical. 

Oxygen  Compound 

Eeducing  agent  taken  up  formed 

Hydrogen  (2  atoms),  H2  +       0  -+  II20 

Carbon  (1  atom),  C  +       0  -»  CO 

Carbon  monoxid  (1  mol.),  CO  +       0  — »  C02 

Sulfur  (1  atom),  S  +     20  ->  S02 

Sulfur  dioxid  (1  mol.),  S02  +       0  ->  S03 

Nitrous  acid  (1  mol.),  HN02  +       0  ->  HN03 

Hydrocyanic  acid   (1  mol.),  HCN  +       0  -*  HCNO 

Stannous  oxid  (1  mol.),  SnO  -|~       O  ~*  SnO2 

Ferrous  oxid  (2  mols.),  2FeO  +       O  -»  Fe203 

The  value  as  reducing  agents  of  the  other  substances  mentioned  may 
be  more  clearly  recognized  by  considering  their  decomposition  products 
and  the  way  in  which  they  oxidize.  Substances  containing  carbon,  hy- 
drogen and  oxygen  usually  oxidize  to  water  and  carbon  dioxid.  Hence 
an  amount  of  oxygen  sufficient  to  convert  any  carbonaceous  substance 
into  H20  and  C02  may  be  taken  as  the  amount  of  oxygen  each  will 
absorb. 

Oxalic  acid,  H2C204,  will  decompose  into  H20,  C02  and  CO.  All  of 
these  compounds  are  oxidized  as  completely  as  possible  with  the  exception 
of  CO.  One  molecule  of  oxalic  acid  is,  therefore,  equivalent  to  ICO. 

Tartaric  acid,  H6C406,  may  break  down  into  3H20  +  SCO  +  C.  One 
molecule  is,  therefore,  equivalent  to  SCO  +  C. 

One  molecule  of  alcohol,  H6C2O,  is  equivalent  as  a  reducing  agent 
to  2H,  since  after  two  hydrogen  atoms  have  been  oxidized  away  from 
each  molecule,  the  residue,  H4C20,  aldehyde,  is  volatile  and  is  removed 
from  the  sphere  of  action. 

Hydrosulfuric  acid,  H2S,  may  be  broken  down  into  2H  and  S.    When 


48  QUALITATIVE    ANALYSIS. 

oxidized  by  permanganate,  the  sulfur  is  largely  converted  into  S03,  but 
with  other  oxidizing  agents  the  main  reaction  is  the  oxidation  of  the 
hydrogen,  the  sulfur  not  being  attacked  readily. 

Sulfurous  acid,  H2SO3,  decomposes  into  H20  and  the  acid  anhydrid, 
S02. 

Thiosulfuric  acid,  H2S203,  will  decompose  into  H2S03  +  S.  Both 
the  S02  and  the  S  are  readily  oxidized  to  S03.  For  the  oxidizing  action 
of  iodin  see  page  189. 

Thiocyanic  acid,  HSCN,  is  equivalent  to  HCN  +  S,  the  sulfur 
being  readily  oxidized  to  S03. 

Hydriodic  acid,  HI,  will  decompose  into  H  + 1.  Iodin  is  not  ordi- 
narily oxidized,  since  it  is  insoluble  in  water  and  volatile  at  100°. 
Hydriodic  acid  is  equivalent  to  1H,  therefore. 

Similarly  Hydrobromic  acid,  HBr,  and  Hydrochloric  acid,  HC1, 
are  each  equivalent  to  1H  for  each  molecule  of  acid  used. 

These  facts  may  be  better  appreciated  when  arranged  in  the  follow- 
ing table  of  equivalents : 

Hydriodic  acid,  HI,  is  equivalent  to  H 

Hydrobromic  acid,  HBr,  H 

Hydrochloric  acid,  HC1,  "  H 

Alcohol,  H6C20,  "  2H 

Hydrosulfuric  acid,  H2S,  "  2H  +  S 

Sulfurous  acid,  H2S03,  "  S02 

Thiosulfuric  acid,  H2S203,  "  S02  +  S 

Thiocyanic  acid,  HSCN,  "  HCN  +  S 

Oxalic  acid,  H2C204,  "  CO 

Tartaric  acid,  H6C4O6,  "  SCO  +  C 

From  the  table  of  oxygen  equivalents  of  the  various  oxidizing  agents 
it  is  seen  that  several  different  substances  are  of  the  same  value,  molecule 
for  molecule.  Two  atoms  of  chlorin  have  the  same  oxidizing  power 
as  two  atoms  of  iodin,  since  each  is  equivalent  to  one  atom  of  oxygen. 
It  is  not  to  be  inferred  from  this,  however,  that  one  substance  gives  as 
general  satisfaction  as  an  oxidizing  agent  as  the  other.  Iodin  is  an 
inert  element  compared  with  chlorin,  and  while  the  latter,  because  of 
its  greater  chemical  activity,  will  attack  and  oxidize  almost  any  sub- 
stance, iodin  has  a  very  limited  application.  Chlorin  is,  therefore,  a 
much  better  oxidizing  agent  for  general  use  than  iodin. 

Similarly  hydrochloric  and  hydriodic  acids  have  the  same  values  as 
reducing  agents,  one  molecule  of  each  being  equivalent  to  one  atom  of 


OXIDATION    AND    KEDUCTION.  49 

hydrogen.  Hydrochloric  acid,  however,  is  a  very  stable  compound  and 
does  not  part  with  its  hydrogen  readily.  Except  in  the  presence  of  the 
very  strongest  oxidizing  agents,  hydrochloric  acid  shows  no  reducing 
properties  at  all.  Hydriodic  acid,  on  the  other  hand,  gives  up  its  hy- 
drogen with  especially  great  readiness.  Hence  hydriodic  acid  is  a  much 
better  reducing  agent  for  ordinary  purposes  than  hydrochloric  acid. 

Much  of  the  value  of  oxidizing  and  reducing  agents  depends  on 
factors  other  than  their  oxygen  or  hydrogen  equivalents.  Chief  among 
these  factors  are  the  chemical  activity  and  stability  of  the  substances. 
Bodies  which  readily  decompose  into  active  decomposition  products  will 
be  the  best  oxidizing  and  reducing  agents.  Substances  of  this  character 
are  known  as  strong  reagents.  Nitric,  dichromic,  permanganic  and 
chloric  acids  together  with  chlorin  and  bromin  are  strong  oxidizing  agents 
while  iodin  is  especially  weak.  Nascent  hydrogen,  hydrogen  sulfid, 
hydriodic  and  sulfurous  acids  are  strong  reducing  agents  while  hydro- 
bromic  acid  and  alcohol  are  weak. 

By  reference  to  the  tables  of  values  already  given,  the  relative 
amounts  of  oxidizing  and  reducing  agents  which  will  interact  may  be 
determined.  As  an  illustration  of  their  use,  let  the  equation  for  the 
oxidation  of  oxalic  acid  by  permanganate  in  acid  solution  be  required. 
Two  molecules  of  permanganic  acid  are  equivalent  to  five  atoms  of 
oxygen.  Each  molecule  of  oxalic  acid  is  equivalent  to  one  molecule  of 
carbon  monoxid  which  will  absorb  one  atom  of  oxygen  and  become 
carbon  dioxid.  Therefore,  five  molecules  of  oxalic  acid  will  be  oxidized 
by  two  molecules  of  permanganic  acid.  The  essential  reaction  is: 


5(70  +  Mn207  -»  5002  +  2MnO 

The  reaction  may  be  completed  by  further  consultation  of  the  tables 
and  the  application  of  the  general  principles  of  chemistry.  Instead  of 
the  equivalent  CO,  H2C2O4  will  appear  on  the  left  side  which  will  intro- 
duce on  the  right  side  1H2O  +  1C02  additional  for  each  molecule  of  ox- 
alic acid  which  is  oxidized.  In  the  presence  of  acid  MnO  will  dissolve 
and  form  a  salt,  hence  if  H2S04  be  the  acid  present,  enough  of  it  must 
appear  in  the  reaction  to  unite  with  the  MnO.  The  complete  reaction 
becomes,  then  : 

5H2C2O4  +  2HMn04  +  2H2S04  ->  10C02  +  8H2O  +  2MnS04 

If  potassium  permanganate  is  used  instead  of  permanganic  acid,  as  is 
usually  the  case,  additional  sulfuric  acid  must  appear  in  the  equation 
5 


50  QUALITATIVE    ANALYSIS. 

sufficient  to  unite  with  the  potassium  oxid  present,  for  2KMnO4  is  equiva- 
lent to  K2O  +  Mn207.  One  molecule  of  sulfuric  acid  will  be  needed  for 
this  purpose  which  will  form  potassium  sulfate  and  water.  Using  potas- 
sium permanganate  the  complete  reaction  becomes: 

5H2C2O4  +  2KMnO4  +  3H2S04  ->  10C02  +  8H20  +  2MnS04  +  K2S04 

The  interaction  between  dichromic  acid  and  hydrogen  sulfid  may  be 
readily  expressed  by  an  equation  when  it  is  known  that  dichromic  acid 
does  not  carry  the  oxidation  any  farther  than  water  and  free  sulfur. 
H2S  is  equivalent  to  2H  +  S.  If  S  is  not  oxidized  further,  H2S  becomes 
equivalent  to  2H.  Two  molecules  of  chromic  anhydrid  are  equivalent 
to  one  molecule  of  the  basic  chromium  trioxid,  Cr203,  +  30.  Since  2H 
unite  with  10,  the  essential  reaction  is  : 

2Cr03  +  3H2S  ->  Cr203  +  3H20  +  3S 

If  the  reaction  takes  place  in  a  solution  of  dilute  hydrochloric  acid,  as  is 
usually  the  case,  the  complete  reaction  becomes  : 

H2O2O7  +  3H2S  +  6HC1  -»  2CrCl3  +  7H20  +  3S 

One  of  the  common  reactions  of  qualitative  analysis  is  the  oxidation 
of  arsenic  sulfid,  As  S:,  with  nitric  acid.  Since  this  substance  is  not 
soluble  in  water  or  acids,  it  is  necessary  to  oxidize  it  to  get  it  into  solu- 
tion. The  reaction, 

As2O5  +  5H2S  -»  As2S5  +  5H2O 

shows  arsenic  sulfid  to  be  a  salt  of  hydrogen  sulfid.  In  this  compound 
arsenic  shows  its  highest  valence  and  can  not  be  oxidized  further.  The 
oxidation  of  As2S5  is,  therefore,  the  oxidation  of  H2S  alone.  Under  the 
conditions  of  the  experiment  much  of  the  sulfur  remains  unoxidized; 
hence  H2S  is  here  equivalent  as  a  reducing  agent  to  2H.  5H2S  will 
require  50.  N2O5  is  equivalent  to  30.  The  essential  equation  is, 
therefore  : 


5N205  +(3X5  =  15)H2S->  5N202  +  15S  +  15H20 
Completing  the  reaction,  we  have  : 


(5  X  2  =  10)HN03  +  3AsA  +  (3  X  5  =  15)H2O'-» 

5N2O2  +  15S  +3As2O5  +  20H2O 

Since,  however,  15H2O  or  more  appears  on  both  sides  of  the  equation, 


OXIDATION    AND    SEDUCTION.  51 

this  quantity  may  be  subtracted  from  each  side  and  the  equation  sim- 
plified as  follows  : 

10HNO8  +  3As2S5  ->  5N202  +  15S  +  3As2O5  +  5H2O 

Many  of  the  oxidation  and  reduction  reactions  applied  in  qualita- 
tive analysis  take  place  in  water  solutions,  hence  are  ionic  reactions. 
Ionic  equations  may  be  readily  written  to  express  them.  For  the  inter- 
action of  oxalic  acid  with  permanganate  : 


2Mn04  +  (5  X  2  =  10)H  +  5CO4  +  (3  X2  = 

2Mn  +  2K  +  3S04  +  10<7O2  +  8H2O 

By  collecting  common  terms  and  eliminating  common  factors  from  both 
sides,  the  equation  becomes  the  simplest  expression  for  the  reaction. 

2Mn04  +  5C  A  +  16H  ->  2Mn  +  10C02  +  8H20 
For  dichromic  and  hydrosulfuric  acids  the  equation  is  : 

2H  +  CrA  +  (  3  X  2  =  6)  H  +  3S  +  6H  +  6C1  -» 

2Cr  +  6C1  +  3S  +  7H20 
The  simplified  expression  is: 

CrA  +  3S  +  1411  -»  2Cr  +  3S  +  7H20 
For  nitric  acid  and  arsenic  pentasulfid  the  equation  is  : 
10H  +  10N03  +  3As2S5  ->  5^202  +  15S  +  5H20 

All  of  these  equations  show  the  disappearance  of  hydrogen  ions 
during  oxidation  reactions.  The  following  ionic  equations  show  the 
decrease  in  hydroxyl  ions  during  reduction: 


_ 
2Fe  +  H2S  +  2  OH->  2Fe  +  S  +  2H20 

In  qualitative  analysis  these  elements  are  not  ordinarily  reduced 
in  alkaline  solution  but  rather  in  acid.  Under  these  conditions  reduction 
causes  the  production  of  hydrogen  ions  which  may  be  considered  equiva- 
lent to  a  decrease  in  hydroxyl  ions  since  on  neutralization  each  disap- 


52  QUALITATIVE    ANALYSIS. 

pears.     The  following  equations  indicate  the  nature  of  the  reduction 
process  in  acid  solutions: 


Striking  from  each  side  of  the  equation  the  common  term  4C1,  the 
expression  becomes: 

§n++2H-      -»     S 


Ferric  iron  reduced  by  hydrogen  sulfid  in  acid  solution  may  be  repre- 
sented as  follows: 

"  2H      S 


These  equations  indicate  clearly  that  in  solutions  of  salts  of  the 
metals  the  essential  feature  of  oxidation  is  an  increase  in  the  positive 
charge  on  the  kathions.  Reduction  correspondingly  reduces  this  charge. 
The  only  difference  between  iron  in  the  ferric  state  and  iron  in  the 
ferrous  state  is  that  ferric  ions  have  three  unit  charges  while  ferrous 
ions  carry  only  two  charges.  The  difference  between  a  charge  of  four 
plus  units  and  two  plus  units  is  the  difference  between  stannic  and 
stannous  ions. 

The  interaction  of  oxidizing  and  reducing  agents  may  be  made  to 
take  place  in  the  "  dry  way  "  by  the  application  of  heat.  Reducing  sub- 
stances, such  as  tartaric  and  oxalic  acids  and  other  carbonaceous  matter 
which  interferes  with  the  systematic  analysis,  are  usually  oxidized  and 
burned  up  completely  by  fusion  with  ammonium,  potassium  or  other  salt 
of  nitric  acid.  Other  oxidizable  substances,  if  present,  are  oxidized  as 
well.  Sulfids  are  changed  to  sulfates,  chromium  compounds  to  chro- 
mates  and  manganese  salts  to  the  green  manganates.  (See  page  263.) 

Many  substances  may  be  oxidized  by  heating  in  the  air,  especially  if 
heated  in  the  flame  of  a  blowpipe.  If  ordinary  illuminating  gas,  con- 
sisting essentially  of  hydrogen  and  carbon  combined  in  various  ways, 
be  mixed  with  air  as  it  emerges  from  the  burner,  complete  combustion 
to  carbon  dioxid  and  water  will  ensue  and  the  flame  will  be  colorless.  The 
flame  from  a  blowpipe  or  Bunsen  burner  is  of  this  character.  Such  a 
flame  consists  of  three  zones.  An  inner  cone  of  unburned  gas  mixed 
with  air  is  separated  from  the  distinctly  visible  intermediate  zone  of 
blue  flame  by  the  surface  where  ignition  begins.  The  hydrogen  con- 
stituent of  the  gas  burns  immediately  to  water.  Carbon  burns  first  to 
carbon  monoxid  which  at  high  temperatures  is  a  strong  reducing 
agent.  Because  of  the  presence  of  this  substance,  many  oxids  held  in 


PRELIMINARY    EXAMINATION.  53 

this  intermediate  zone  are  reduced,  hence  this  portion  is  known  as  the 
reducing  flame.  Just  beyond  the  tip  of  the  distinct  blue  cone  the  tem- 
perature is  highest  and  the  best  reduction  effects  are  obtained  here. 

Carbon  monoxid  burns  to  carbon  dioxid  with  a  flame  of  character- 
istic blue  color.  The  products  of  combustion,  water  and  carbon  dioxid, 
constitute  the  outer,  scarcely  visible  zone  of  the  flame.  While  still  very 
hot  these  gases  become  mixed  with  oxygen  from  the  air.  Since  oxygen 
at  high  temperatures  is  a  very  active  agent,  oxidation  takes  place  very 
readily  in  this  outer  zone  which  is  known  as  the  oxidizing  flame.  Just 
beyond  the  point  of  the  visible  flame  oxidation  takes  place  most  rapidly. 

By  heating  a  small  particle  of  a  substance  before  the  blowpipe  in 
the  proper  position  in  the  flame,  either  oxidation  or  reduction  may  be 
effected.  To  obtain  a  satisfactory  blowpipe  flame  the  supply  of  air 
and  gas  entering  a  Bunsen  burner  should  be  cut  down  until  a  luminous 
flame  not  more  than  one  inch  high  is  obtained.  By  directing  the  blow- 
pipe across  the  top  of  the  burner  and  blowing  gently  yet  steadily,  a  flame 
may  be  obtained  in  which  the  zones  will  be  definitely  separated  and  sat- 
isfactory results  obtained. 

PRELIMINARY  EXAMINATION. 

By  the  application  of  heat  many  solids  exhibit  phenomena  so  char- 
acteristic as  to  furnish  evidence  as  to  their  nature.  Consequently  a 
series  of  very  simple  tests  which  may  be  quickly  performed  has  been 
arranged  in  the  hope  that,  by  such  a  preliminary  examination,  informa- 
tion of  value  in  the  systematic  analysis  may  be  obtained.  If  the  sub- 
stance be  homogeneous  and  fairly  simple,  a  preliminary  examination  alone 
may  be  sufficient  to  determine  its  identity.  The  methods  ordinarily  used 
in  making  a  preliminary  examination  are:  (a)  heating  in  a  closed-tube, 
(6)  heating  with  concentrated  sulfuric  acid,  (c)  heating  on  charcoal,  (d) 
heating  in  a  borax  bead,  (e)  heating  on  a  platinum  wire  in  a  Bunsen 
flame. 

(a)  Heating  in  a  closed-tube  consists  in  heating  over  a  Bunsen  burner 
a  small  amount  of  the  substance  under  examination  in  a  small  tube  of 
hard  glass  closed  at  one  end.  The  object  of  the  tube  is  to  prevent  as  far 
as  possible  the  oxidizing  influence  of  the  air  in  order  that  a  reaction  due 
to  increase  of  temperature  alone  may  be  observed. 

Under  these  conditions  easily  fusible  substances  melt  and  may  often! 
be  volatilized  completely,  sometimes  giving  off  very  characteristic  odors. 
Other  substances  do  not  melt  under  ordinary  atmospheric  pressure,  but 
pass  immediately  into  the  gaseous  state.  Such  substances  are  said  to 


54  QUALITATIVE    ANALYSIS. 

sublime.  On  cooling,  the  vapors  condense  to  the  solid  form  without  pass- 
ing through  the  intermediate  liquid  condition.  Where  it  may  be  satis- 
factorily applied,  sublimation  offers  an  easy  method  of  purification  from 
non-volatile  impurities. 

Many  difficultly  fusible  substances  suffer  a  change  in  color  on  being 
heated,  which  in  some  cases  is  very  characteristic. 

Many  substances  suffer  decomposition  by  heating  before  they  become 
volatile  or  in  some  cases  even  before  melting.  Reactions  in  which  certain 
of  the  decomposition  products  are  volatile  under  the  conditions  of  the 
experiment  are  most  likely  to  take  place,  since  a  reversal  of  the  reaction 
is  prevented  by  the  removal  of  the  volatile  constituents  as  fast  as  they 
are  formed.  Water,  being  very  volatile,  is  one  of  the  commonest  decom- 
position products  of  high  temperature  reactions. 

When  heated  sufficiently  high,  all  bases  lose  water  and  form  oxids 
of  the  metals  which  are  generally  quite  stable  and  are  not  further 
decomposed. 

Acids  decompose  into  water  and  the  anhydrid  generally.  Some 
anhydrids  are  very  stable.  Others  are  readily  decomposed  and  break 
down  still  further. 

As  a  rule,  salts  decompose  into  acid  anhydrids  and  oxids  of  the 
metals. 

In  general,  the  more  stable  the  acid  is,  the  more  stable  will  the 
salts  be.  Hypochlorous,  nitrous,  carbonic  and  sulfurous  acids  are  all 
unstable  at  ordinary  temperatures.  At  higher  temperatures  their  salts 
are  correspondingly  unstable.  Nitric,  chloric  and  sulfuric  acids  are 
more  stable  than  the  last  mentioned  acids  and  their  salts  require  corre- 
spondingly higher  temperatures  to  decompose  them.  Finally,  boric  and 
phosphoric  acids  are  very  stable  bodies  and  their  salts  are  more  resistant 
than  others. 

The  salts  of  a  given  acid  with  the  various  metals  possess  varying  de- 
grees of  stability.  The  most  stable  salts  are  formed  by  the  most  strongly 
basic  elements.  To  illustrate,  the  carbonates  of  ferric  iron,  chromium, 
tin,  arsenic  and  antimony  are  not  known  since  they  are  unstable  at 
ordinary  temperatures.  The  carbonates  of  mercury,  copper,  lead,  mag- 
nesium and  calcium  show  increasing  stability,  but  they  are  all  decom- 
posed at  a  dull  red  heat.  Strontium  and  barium  carbonates  require  the 
highest  temperature  of  the  Bunsen  flame  while  the  carbonates  of  the 
alkali  metals  are  undecomposed  even  at  this  temperature. 

None  of  the  oxids  of  the  metals  are  particularly  volatile,  hence  am- 
monia is  the  only  basic  salt-forming  substance  driven  off  by  the  heat  of 


PKELIMINAKY   EXAMINATION.  55 

a  Bunsen  burner.*  Most  of  the  acid  anhydrids  are  volatile  at  this 
temperature,  silicic,  boric  and  phosphoric  anhydrids  being  the  excep- 
tions. Therefore,  silicates,  borates  and  phosphates  generally  are  among 
the  most  stable  substances  at  high  temperatures,  since  neither  acid  nor 
basic  oxid  is  volatile. 

When  an  acid  anhydrid  is  unstable  and  decomposes  further,  the  de- 
composition temperatures  of  its  salts  do  not  show  the  wide  variation 
noticed  with  the  carbonates  where  the  acid  anhydrid  is  volatile  only.  All 
nitrates  decompose  readily  at  about  the  same  temperature  into  a  metallic 
oxid  and  the  acid  anhydrid,  N,05.  The  latter  undergoes  further  decom- 
position in  the  manner  set  forth  on  page  150. 

Fluorids,  chlorids,  bromids  and  iodids  are  very  stable,  although  some 
salts  are  quite  volatile. 

The  decomposition  of  the  salts  of  cyanogen  acids  depends  consider- 
ably on  the  nature  of  the  metallic  element  present  and  will  be  treated 
in  detail  under  appropriate  headings  when  necessary. 

Non-volatile  organic  salts  and  acids  generally  suffer  far-reaching 
decompositions  when  heated,  resulting  in  the  separation  of  more  or  less 
free  carbon.  This  process  is  known  as  charring  or  carbonization. 

Certain  of  these  phenomena  are  often  observed  by  heating  in  a  closed- 
tube.  The  characteristic  changes  and  their  significations  are  given 
on  pages  243  and  244. 

( b )  By  heating  with  concentrated  sulf uric  acid  in  a  test-tube,  decom- 
positions analogous  to  those  considered  are  brought  about.     Such  reac- 
tions take  place  more  rapidly  and  at  lower  temperatures  in  the  presence 
of  the  acid  than  without  this  reagent.     Decomposition  products  are  fre- 
quently liberated  which  are  easily  recognized  and  afford  indications  as  to 
the  composition  of  the  substances  so  treated.     For  characteristic   de- 
compositions see  pages  245. 

(c)  Heating  on  charcoal  may  be  any  one  of  three  processes,  depending 
on  the  nature  of  the  substance  and  the  conditions  of  the  experiments. 
(1)  If  the  substance  be  not  reactive  at  high  temperatures,  any  change 
observed  will  probably  be  physical.     Fusion,  volatilization,   change  of 
color,  formation  of  a  sublimate  condensing  on  the  charcoal  at  a  distance 
from  the  flame,  all  may  be  observed.     (2)  If  the  charcoal  be  simply  used 
as  a  support,  the  substance  may  be  oxidized.      (3)   If  the  substance  is 
mixed  with  sodium  carbonate,  used  as  a  flux  and  to  protect  from  the  action 
of  the  air,  the  charcoal  may  exert  a  reducing  action,  especially  if  some  of 

*  Mercuric  oxid  may  be  decomposed  and  then  volatilized  by  the  heat  of  a  Bunsen 
flame. 


56  QUALITATIVE    ANALYSIS. 

it  is  finely  powdered  and  mixed  with  the  sodium  carbonate.  From  basic 
oxids  certain  of  the  metals  may  be  separated  in  this  way.  Characteristic 
phenomena  and  their  indications  may  be  found  on  pages  246  and  247. 

(d)  Heating  in  a  borax-bead  before  the  blowpipe  is  a  high  tempera- 
ture reaction  frequently  productive  of  very  characteristic  colors.     On 
being  heated,   borax,  Na2B407.10H20,   intumesces   and  loses   its  water 
of  crystallization.     It  then  fuses  to  a  clear,  colorless  glass  which  has  the 
power  of  dissolving  metallic  oxids.    Borax  may  be  considered  as  a  com- 
pound of  two  molecules  of  sodium  metaborate,  NaB02,  with  one  molecule 
of  boric  acid  anhydrid,  B203,  the  combination  being  so  unstable  under 
certain  conditions  that  it  readily  breaks  down  into  these  components. 
There  is,  therefore,  such  a  predominance  of  the  acicT  constituent  that 
when  borax  is  heated  with  metallic  oxids,  the  acid  anhydrid  unites  with 
them  to  form  metaborates  of  the  metals.     The  reaction  with  cobalt  is 
typical. 

Na2B407  +  CoO  -»  2NaB02  +  Co(B02)2 

The  metaborates  of  certain  of  the  heavy  metals  are  very  intensely  col- 
ored, a  mere  trace  of  the  oxid,  or  of  a  substance  capable  of  breaking 
down  into  the  oxid  under  the  conditions  of  the  experiment,  being  dis- 
tinctly indicated  by  the  borax  bead  test.  The  characteristic  colors  de- 
veloped are  given  on  page  247. 

(e)  The  heat  of  the  Bunsen  flame  is  sufficient  to  volatilize  the  com- 
pounds of  certain  metals  and  to  cause  momentary  decomposition,  in  some 
cases  into  the  free  metal.     In  this  condition,  many  salts  produce  a  dis- 
tinct coloration  -  of  the  Bunsen  flame  which  is  characteristic  of  the 
metallic  component.     (See  page  247.) 

Of  the  common  salts  chlorids  dissociate  most  readily,  hence  it  is 
advisable  to  moisten  with  hydrochloric  acid  a  small  portion  of  the 
finely  powdered  substance.  The  paste  may  be  inserted  in  a  small  loop  of 
platinum  wire  and  heated  just  beyond  the  blue  cone  in  a  Bunsen  flame. 

The  coloration  produced  by  certain  metals  is  so  intense  that  the 
merest  trace  of  their  salts  is  sufficient  to  indicate  their  presence.  Hence 
it  is  necessary  that  the  platinum  wire  should  be  very  clean.  A  satis- 
factory method  of  cleaning  a  platinum  wire  from  anything  which  may 
contaminate  it  is  to  make  a  small  borax  bead  and  cause  it  to  pass  along 
the  wire  by  holding  it  in  such  a  position  in  the  flame  that  the  bead  is 
heated  on  one  side  only.  Under  these  conditions  the  bead  will  travel 
away  from  the  flame  along  the  wire  and  dissolve  anything  that  may  be 
present.  By  driving  the  bead  to  the  end  of  the  wire  and  heating  strongly 


WEIGHTS   AND    MEASUEES.  57 

for  a  moment,  the  bead  may  be  thrown  off  by  a  sharp  jerk,  leaving  the 
wire  free  from  everything  except  a  small  amount  of  borax.  This  may 
be  removed  by  alternately  dipping  the  wire  in  pure,  dilute  hydrochloric 
acid  and  heating  until  the  sodium  color  is  no  longer  evident. 

The  color  produced  by  insignificant  amounts  of  sodium  is  so 
intense  as  to  mask  very  completely  the  indication  given  by  relatively  large 
amounts  of  potassium.  For  this  reason  the  detection  of  potassium  in  the 
presence  of  sodium  becomes  difficult  to  the  unaided  eye.  By  the  spectro- 
scope each  may  be  detected  in  the  presence  of  the  other,  since  each  is  indi- 
cated by  separate  and  distinct  lines.  Without  the  spectroscope  it  is 
necessary  to  cut  out  the  sodium  indication  by  something  which  will 
absorb  the  yellow  rays  of  light  produced  by  it  and  at  the  same  time  trans- 
mit the  blue  rays  produced  by  potassium. 

The  most  satisfactory  method  of  accomplishing  this  point  is  by 
the  use  of  a  bottle  filled  with  blue  dye  of  such  strength  that  a  strong 
sodium  flame  is  invisible  when  viewed  through  this  solution.  This 
apparatus  is  known  as  a  potassioscope.  With  its  aid  traces  of  potassium 
may  be  detected  in  the  presence  of  quantities  of  sodium.  By  suitably 
choosing  the  dye,  not  only  the  sodium  flame,  but  also  the  indication  given 
by  all  other  common  elements  except  potassium  may  be  absorbed  and 
thus  all  possible  confusion  avoided.  (See  page  319.) 

WEIGHTS  AND  MEASUEES 

The  atoms  and  molecules  which  enter  into  and  constitute  chemical 
substances  have  different  relative  weights  which  may  be  ascertained  read- 
ily. The  weight  of  the  atom  of  hydrogen  is  taken  as  the  unit  for  com- 
parison and  the  atomic  weight  of  hydrogen  is  said  to  be  1.  That  num- 
ber is  taken  as  the  weight  of  any  other  atom  which  tells  how  many  times 
heavier  it  is  than  the  hydrogen  atom;  thus  oxygen  is  16,  sodium  23,  sulfur 
32  and  chlorin  35.5.  The  atomic  weights  of  the  elements  are  given  on 
page  324. 

The  weight  of  the  molecule  is  equal  to  the  sum  of  the  weights  of  the 
atoms  in  it ;  thus  the  molecular  weight  of  NaCl  is  58.5 ;  Na2S04  is  142. 

When  one  gram  of  any  substance  is  taken  for  each  unit  in  its 
molecular  weight,  this  amount  is  called  a  gram-molecule.  Thus  40  grams 
of  NaOH  constitutes  a  gram-molecule,  for,  since  the  atomic  weight 
of  sodium  is  23,  of  oxygen  16,  and  of  hydrogen  1,  there  are  40  units  in 
the  molecular  weight.  In  the  gram-molecule  of  NaOH  there  are  23  grams 
of  sodium,  16  grams  of  oxygen  and  1  gram  of  hydrogen.  Since  the 
molecular  weight  of  H2SO4  is  98,  in  the  gram-molecule  of  sulfuric  acid 


58  QUALITATIVE    ANALYSIS. 

there  are  32  grams  of  sulfur,  64  grams  of  oxygen  and  2  grams  of  hydro- 
gen, making  98  grams  in  all. 

Solutions  containing  a  gram-molecule  of  a  substance  in  a  liter  of 
solution  are  called  gram-molecular  solutions.  In  a  gram-molecule  of 
all  substances  there  is  the  same  number  of  molecules  and  in  equal  volumes 
of  gram-molecular  solutions  there  is  the  same  number  of  molecules. 

A  molecule  of  one  substance  reacts  usually  with  a  single  molecule 
of  another  substance  or  with  some  simple  multiple  of  this  proportion,  as 
1 :  2,  1 :  3,  1 :  4,  2  :  3,  and  similar  ratios.  Consequently  one  gram-molecule 
of  one  substance  reacts  with  a  single  gram-molecule  of  another  substance 
or  with  some  simple  multiple  of  this  ratio.  Thus  1  gram-molecule  of  the 
monoacid  base,  NaOH,  is  just  sufficient  to  neutralize  1  gram-molecule  of 
the  monobasic  acid,  HC1.  One  gram-molecule  of  the  dibasic  acid,  H2S04, 
will  require  2  gram-molecules  of  the  monoacid  base,  NaOH,  to  neutralize 
it  or  1  gram-molecule  of  the  diacid  base,  Ca(OH)2. 

In  a  gram-molecule  of  any  monobasic  acid  there  is  one  gram  of 
hydrogen  capable  of  passing  into  the  ionic  condition  and  of  being  re- 
placed by  metal  in  the  formation  of  salts.  In  a  gram-molecule  of  any 
dibasic  acid  there  are  two  grams  of  this  hydrogen  while  a  tribasic  or 
tetrabasic  acid  contains  three  or  four  grams.  One-half  a  gram-molecule 
of  any  dibasic  acid  contains  as  much  hydrogen  capable  of  passing  into 
the  ionic  condition  as  a  gram-molecule  of  any  monobasic  acid  and  is 
therefore  equivalent  to  it  as  an  acid.  Similarly  one-third  of  a  gram- 
molecule  of  a  tribasic  acid  and  one-quarter  of  a  gram-molecule  of  a 
tetrabasic  acid  are  equivalent  to  a  gram-molecule  of  any  monobasic  acid. 

One  gram-molecule  of  a  monoacid  base  is  just  sufficient  to  completely 
neutralize  one  gram-molecule  of  a  monobasic  acid,  hence  the  two  quanti- 
ties may  be  said  to  be  equivalent.  One  equivalent  of  salt  is  formed  by 
the  union  of  one  equivalent  of  acid  and  of  base.  One-half  of  a  gram- 
molecule  of  a  diacid  base,  one-third  a  gram-molecule  of  a  triacid  base 
and  one-fourth  of  a  gram-molecule  of  a  tetracid  base  are  equivalent  to 
one  gram-molecule  of  a  monoacid  base. 

In  this  way  a  system  of  equivalents  has  been  developed  such  that 
an  equivalent  of  any  substance  is  just  sufficient  to  react  completely  with 
an  equivalent  of  any  other  substance.  That  amount  is  taken  as  the 
equivalent  of  any  substance  which  will  combine  with  or  take  the  place 
of  one  gram  of  hydrogen.  Thus  35.5  grams  of  chlorin  unite  with  1  gram 
of  hydrogen  to  form  36.5  grams  of  hydrogen  chlorid.  Therefore  35.5 
grams  of  chlorin  and  36.5  grams  of  hydrogen  chlorid  are  the  equivalents 
of  1  gram  of  hydrogen.  35.5  grams  of  chlorin  may  be  replaced  by  17 


WEIGHTS    AND    MEASUEES.  59 

grams  of  hydroxyl,  OH,  to  form  water.  17  grams  of  the  OH  radicle  is 
the  equivalent  of  1  gram  of  hydrogen.  1  gram  of  hydrogen  may  be 
replaced  by  23  grams  of  sodium  or  39  grams  of  potassium  or  108  grams 
of  silver,  hence  these  amounts  are  all  equivalents. 

40  grams  of  calcium  unite  with  71  grams  of  chlorin.  Since  35.5 
grams  of  chlorin  are  equivalent  to  1  gram  of  hydrogen,  71  grams  are  2 
equivalents.  Then  40  grams  of  calcium  is  2  equivalents  also,  and  20 
grams  of  calcium  is  the  equivalent  of  1  gram  of  hydrogen. 

In  this  way  the  values  of  the  equivalents  of  each  element  and  com- 
pound may  be  determined.  In  general,  the  equivalent  of  an  acid  or  base 
equals  the  gram-molecule  divided  by  the  basicity  or  acidity  of  the 
compound.  The  equivalent  of  an  element  is  equal  to  the  atomic  weight 
divided  by  the  valence.  The  equivalent  of  any  compound  is  that  part 
of  a  gram-molecule  which  contains  one  equivalent  of  the  active  agent. 

By  the  term  active  agent  is  understood  that  particular  constituent  of 
the  compound  which  is  the  essential  factor  in  the  reaction  under  con- 
sideration. A  compound  containing  several  different  elements  in  differ- 
ent proportions  may  be  considered  from  the  standpoint  of  each  element 
in  different  reactions,  hence  its  equivalent  may  vary.  Potassium  per- 
manganate, KMn04,  may  act  as  a  potassium  salt.  Under  these  con- 
ditions potassium  is  the  essential  factor.  39  grams  of  potassium  are 
contained  in  158  grams  of  permanganate  and  the  gram-molecule  would 
be  the  equivalent  under  these  circumstances.  Again,  the  compound  may 
be  used  as  a  source  for  the  divalent  Mn  ion.  Manganese  has  an  atomic 
weight  of  55,  hence  its  equivalent  is  27.5  grams.  Since  the  equivalent 
of  manganese  is  contained  in  79  grams  of  permanganate,  the  equivalent 
of  permanganate  regarded  as  a  source  for  Mn  ions  is  one-half  the  gram- 
molecule.  On  the  other  hand,  from  one  gram-molecule  of  KMn04  40 
grams  of  oxygen  are  readily  liberated.  Since  8  grams  of  oxygen  unite 
with  1  gram  of  hydrogen,  40  grams  of  oxygen  is  5  equivalents.  Con- 
sidering oxygen  as  the  active  agent,  one-fifth  of  a  gram-molecule  of  per- 
manganate is  the  equivalent  of  1  gram  of  hydrogen. 

If  a  solution  be  made  of  such  strength  that  in  a  litre  of  solution  there 
is  one  equivalent  of  solute,  it  is  known  us  a  normal  solution.  Solutions 
of  this  strength  are  designated  by  the  capital  letter  N;  of  twice  this 
strength  by  2N ;  of  ten  times  this  strength  by  ION ;  one-half  or  one-tenth 
this  strength  by  N/2  or  N/10-.  Equal  volumes  of  normal  solutions  are 
just  sufficient  to  react  together  completely.  Thus  1  liter  N  NaOH  or 
KOH  or  Ca(OH)2  will  just  neutralize  1  liter  of  N  H2S04  or  HNO3  or 
HC1.  1  ccm.  N  Nad  solution  will  completely  precipitate  1  ccm.  N 


60  QUALITATIVE    ANALYSIS. 

AgN03  solution.    10  ccm.  2N  Na2C03  is  just  sufficient  to  react  completely 
with  200  ccm.  N/10  CaCl2  solution. 

For  purposes  of  convenience  in  handling,  laboratory  reagents  are 
generally  made  up  according  to  this  method.  The  strengths  of  the 
ordinary  reagents  suitable  for  qualitative  analysis  are  given  on  pages 
313-320,  also  directions  for  preparing  these  solutions. 


PAKT  II 


DESCRIPTIVE 


The  Classification  of  the  Elements. 

A  detailed  study  of  the  elements  and  of  the  principal  compounds 
formed  by  each  reveals  the  fact  that,  while  any  given  element  is  very 
unlike  most  others  in  nature  and  properties,  it  does  show  very  close  and 
striking  similarity  to  a  few  of  them.  If  a  systematic  consideration  of 
the  elements  be  taken  up  in  the  order  of  their  increasing  atomic  weights, 
hydrogen  being  excluded  since  it  does  not  resemble  any  other  elementary 
substance,  we  begin  with  the  rare  element,  helium,  At.  Wt.  4.  This  is 
chiefly  remarkable  for  its  inertness.  No  compounds  of  it  are  known. 
Since  it  does  not  seem  to  have  the  power  to  unite  with  other  elements, 
its  valence  may  be  said  to  be  zero.  The  second  element  is  lithium,  At. 
Wt.  7,  a  typical  metal,  forming  many  compounds  in  which  it  shows  a 
valence  of  one.  Its  hydroxid  acts  as  a  very  strong  base.  The  third 
element  is  beryllium,  At.  Wt.  9,  a  divalent  metal,  uniting  readily  with 
other  elements.  Its  hydroxid  acts  as  a  base,  but  it  is  not  as  strong  as 
lithium  hydroxid.  Trivalent  boron,  with  At.  Wt.  11,  is  the  fourth 
element.  Here  metallic  properties  have  fallen  off  very  materially  while 
the  characteristics  of  non-metallic,  acid-forming  elements  begin  to  ap- 
pear. Tetravalent  carbon,  At.  Wt.  12,  is  the  next  element,  showing 
acid-forming  rather  than  base-forming  properties.  Next  in  order  comes 
nitrogen,  At.  Wt.  14.  It  is  a  typical  non-metallic  element,  forming  com- 
pounds in  which  valences  of  three  or  five  are  shown.  Oxygen,  At.  Wt. 
16,  next  in  order,  possesses  strong  non^metallic  properties  and  exhibits 
a  valence  of  two  (possibly  four  or  six)  in  its  compounds.  Fluorin, 
At.  Wt.  19,  is  the  most  non-metallic  of  all  elements,  forming  compounds 
in  which  a  valence  of  one  is  characteristic. 

Although  the  atomic  weights  of  these  substances  are  not  widely  differ- 
ent, no  one  of  the  elements  can  be  considered  as  resembling  any  of  the 
others,  and,  while  the  properties  of  fluorin  and  lithium  seem  to  be  as 
different  as  is  possible,  yet,  considering  the  whole  series  just  mentioned, 
only  a  gradual  change  from  the  one  to  the  other  is  to  be  noted. 

Continuing  the  list,  the  next  element  is  neon,  At.  Wt.  20,  remarkably 
inert,  forming  no  compounds  and  resembling  helium  so  closely  as  to 
make  a  separation  of  them  exceedingly  difficult.  The  next  member, 
sodium,  At.  Wt.  23,  is  a  typical  univalent  metal.  Its  hydroxid  is  one  of 
the  strongest  bases  and,  in  every  way,  it  resembles  lithium  very  closely. 
Magnesium,  At.  Wt.  24,  is  a  divalent  metal  resembling  beryllium. 


CLASSIFICATION    OF    THE    ELEMENTS.  63 

Aluminum,  At.  Wt.  27,  is  trivalent  and  analogous  to  boron,  while  silicon, 
At.  Wt.  28,  is  tetravalent  and  like  carbon.  Phosphorus,  At.  Wt.  31, 
duplicates  most  of  the  compounds  of  nitrogen.  Sulfur,  At.  Wt.  32,  re- 
sembles oxygen  and  shows  generally  a  valence  of  two  or  six.  Chlorin, 
At.  Wt.  35,  possesses  non-metallic  properties  in  a  very  high  degree,  forms 
several  of  the  strongest  acids  and  shows,  among  others,  characteristic 
valences  of  one  and  seven. 

Considering  the  members  of  what  may  be  termed  this  second  series, 
we  note  similarity  so  close  as  to  seem  almost  a  duplication  of  the  proper- 
ties of  the  first  series.  The  same  change  from  strongly  metallic  to  de- 
cidedly non-metallic  properties,  accomplished  by  a  gradual  gradation 
through  intervening  elements,  is  observed.  No  member  of  this  series  can 
be  said  to  be  like  any  other  member,  yet  each  is  very  like  the  correspon- 
ding member  of  the  first  series. 

Continuing  the  consideration  of  the  elements,  the  same  variation  and 
similarities  are  observed,  for  argon  is  like  neon  and  helium,  potassium 
like  sodium  and  lithium,  calcium  like  magnesium  and  beryllium,  ect. 
In  short,  if  the  elements  be  arranged  in  series  according  to  their  increas- 
ing atomic  weights,  spaces  being  left  in  some  instances  for  hypothetical 
elements  as  yet  undiscovered,  the  elements  will  be  found  to  arrange  them- 
selves in  groups,  each  group  containing  elements  which  possess  similar 
properties.  From  this  it  would  seem  that  the  properties  of  the  elements 
depend  in  some  way  upon  their  atomic  weights.  An  increase  in  atomic 
weight  at  first  produces  entirely  different  properties,  but,  after  a  period, 
gives  rise  to  properties  closely  resembling  those  previously  noted. 

As  first  enunciated  by  Mendeleeff,  this  generalization,  known  as  the 
Periodic  Law,  states  that  the  properties  of  the  elements  are  periodic  func- 
tions of  their  atomic  weights.  That  is,  that  the  properties  of  the  elements 
depend  primarily  on  their  atomic  weights  and  that,  when  arranged  in 
the  order  of  their  increasing  atomic  weights,  elements  of  a  certain  type 
occur  periodically. 

The  Periodic  System  is  a  method  of  classification  arranging  the  ele- 
ments into  groups  such  that  the  members  of  each  group  show  great  chem- 
ical similarity  to  each  other.  At  the  top  of  each  group  as  given  in  the 
chart,  the  principal  valences  of  the  elements  in  the  group  are  shown  by 
typical  formulae,  R  standing  for  any  element  belonging  to  the  group. 
The  increase  in  valence  from  zero  to  four  and  the  subsequent  decrease  in 
valence  toward  hydrogen  is  shown  by  the  formulae  for  hydrogen  com- 
pounds, while  the  steady  increase  in  valence  toward  oxygen  is  indicated 
by  the  formulae  for  oxids. 


64 


PEEIODIC    SYSTEM. 


Series. 

Group  0. 

Group  1. 
RH 
R20 

Group  2. 
RHa 
RO 

Group  3. 
RH3 
R20S 

Group  4. 
RH. 
R02 

1 
2 
3 
4 
5 
6 
7 
8 
9 
10 
11 

He       4 

Ne     20 
Ar     40* 

Kr     82 
Xe   128 

Li      7 
Na   23 
K     39 
63  Cu 
Eb   85 
108  Ag 
Cs  133 

197  Au 

Be       9 
Mg  24 
Ca    40 
65  Zn 
Sr     88 
112  Cd 
Bal37 

200  Hg 
Ba225 

B       11 

Al    27 
Sc    44 
70  Ga 
Yt    89 
115  In 
La  139 

204  Tl 

C        12 

Si     28 
Ti     48 
72  Ge 
Zr     91 
119  Sn 
Ce  140 

207  Pb 
TH232 

Certain  very  rare  substances,  at  present  thought  to  be  elementary  in  character, 
weights,  approximately  determined,  are:  Praseodymium — 140;  Neodymium — 144 j 
Ytterbium — 173.  *  Probably  incorrectly  determined. 

In  a  general  way,  each  one  of  the  groups  of  the  periodic  system  may 
be  said  to  possess  certain  group  properties  which  are  characteristic  of 
each  member  in  the  group.  Chief  among  these  are  the  valence  and  the 
general  character  of  the  elements ;  i.  e.,  do  the  elements  in  any  particular 
group  act  normally  as  strong  acid-formers  or  strong  base-formers,  or  are 
they  of  an  intermediate  nature? 

Within  each  group,  however,  certain  gradations  are  to  be  noted. 
The  most  important  of  these  is  the  fact  that  as  the  atomic  weight  of  the 
elements  in  any  particular  group  increases,  non-metallic  properties  grow 
less,  while  metallic  properties  increase.  This  holds  true  for  all  groups 
whether  they  be  of  a  general  acid-forming,  base-forming  or  neutral 
character.  Thus,  in  the  strong  base-forming  group  caesium,  At.  Wt. 
133,  is  more  metallic  than  lithium,  At.  Wt.  7.  In  the  decidedly  non- 
metallic  halogen  group,  fluorin,  At.  Wt.  19,  shows  this  non-metallic 
character  in  a  much  greater  degree  than  does  iodin,  At.  Wt.  127.  In  the 
intermediate  fifth  group,  nitrogen,  At.  Wt.  14,  is  a  typical  acid-forming 
element  while  bismuth,  At.  Wt.  209,  is  typically  metallic  and  a  weak 
base-forming  element. 

As  the  atomic  weight  of  the  elements  rises  above  fifty,  a  differentiation 
of  the  elements  in  each  group  according  to  two  types  is  to  be  noted,  and, 
while  both  types  generally  possess  the  group  characteristics,  they  clearly 
differ  from  each  other  in  many  ways.  Thus,  in  the  sixth  group,  chro- 
mium, molybdenum,  tungsten  and  uranium  are  quite  different  from 
oxygen,  sulfur,  selenium  and  tellurium,  although  all  these  elements  may 
with  propriety  be  classed  in  one  group. 


PERIODIC    SYSTEM. 


65 


Group  5. 
KH3 
RoOs 

Group  6. 
RH2 
R03 

Group  7. 
RH 
R207 

Group  8. 

Series. 

N     14 

0      16 

F      19 

1 

P     31 

S     32 

Cl    35 

2 

51  Va 

52  Cr 

55  Mn 

Fe    56     Co     59     Ni     59 

3 

As  75 

Se    79 

Br   80 

4 

94  Nb 

96  Mo 



Kul02     Khl03     Pd  106 

5 

Sbl20 

Tel27* 

I    127 

6 









7 

183  Ta 

184  Wo 

_____ 

Os   191     Ir   193     Pt  195 

9 

Bi  209 





10 



238  U 



11 

are  not  placed  in  the  table  because  of  insufficient  data.     Their  names  and  atomic 
Samarium — 150;    Gadolinium — 156;    Terbium — 160;    Erbium — 166;    Thulium — 171; 
*  Probably  incorrectly  determined. 

In  accordance  with  the  classification  afforded  by  the  Periodic  System, 
the  more  commonly-occurring  elements  will  be  considered  in  detail. 

Group  0. 

Helium  4 ;        Neon  20 ;        Argon  39.9 ;        Krypton  81.8 ;        Xenon  128. 

Of  the  substances  included  in  this  group,  argon  only  is  at  all  common. 
It  occurs  in  the  atmosphere  to  the  extent  of  about  one  per  cent,  along 
with  all  the  other  members  of  this  group,  the  latter  being  present  in 
exceedingly  minute  quantities.  All  are  gaseous  and  are  characterized 
chiefly  by  their  lack  of  chemical  activity.  Since  it  has  not  been  found 
possible  to  make  any  of  them  enter  into  combination  with  any  other 
substance,  they  may  be  said  to  have  a  valence  of  zero.  Only  one  type  of 
element  within  this  group  has  been  discovered.  Because  of  their  rarity 
and  lack  of  compounds,  these  elements  will  not  be  considered  in  detail. 

Group  1 ;  First  Type. 

Lithium  7;          Sodium  23;          Potassium  39.2;          Kubidium  85.5; 
Caesium  132.9;          Ammonium. 

This  group  of  alkali  metals,  as  they  are  called,  are  soft,  light,  lustrous 
substances,  possessing  metallic  properties  in  a  marked  degree.  They  are 
exceedingly  reactive,  tarnishing  in  the  air  and  decomposing  water  with 
the  liberation  of  hydrogen  and  the  formation  of  hydroxids  which  are  the 
strongest  of  bases.  Their  compounds  are  generally  white  in  color  and 
very  soluble,  and  the  group  is  to  be  remembered  as  the  only  one  forming 

6 


66  QUALITATIVE    ANALYSIS. 

soluble  carbonates  and  phosphates.  Unlike  all  other  groups,  the  hydrox- 
ids  and  carbonates  of  these  elements  are  not  decomposed  by  heat  and  the 
only  acid  carbonates,  known  in  the  solid  state,  are  those  of  the  alkali 
metals.  So  closely  do  the  different  elements  resemble  each  other  that 
separations  are  oftentimes  difficult.  Heated  in  a  Bunsen  flame,  however, 
their  compounds  produce  very  characteristic  colors  by  which  the  ele- 
ments may  be  detected  and  differentiated. 

Ammonium  hydroxid  is  known  as  the  ''volatile  alkali"  in  distinction 
from  the  other  "fixed  alkalies,"  i.  e.,  non- volatile  alkalies,  just  men- 
tioned. Because  its  compounds  possess  properties  very  similar  to  those 
of  potassium  compounds,  it  will  be  considered  along  with  these  elements. 

Because  of  the  great  solubility  of  the  salts  of  the  alkali  metals,  they 
are  much  used  to  introduce  various  anions  into  solutions  containing 
the  ions  of  other  metals,  for  the  purpose  of  bringing  about  reactions. 

Rubidium  and  caesium  are  very  rare  and  will  not  be  considered 
in  detail. 

LITHIUM.    Li  — 7- 

Lithium  is  a  silver-white  metal.  It  is  the  lightest  of  all  solid  sub- 
stances and  will  float  on  the  surface  of  petroleum  and  most  other  liquids 
which  might  be  used  to  protect  it  from  the  action  of  the  atmosphere.  It 
resembles  sodium  (see  page  67)  so  closely  in  all  of  its  properties  that  a 
detailed  description  is  unnecessary. 

-  Lithium  salts  are  generally  white  and,  with  the  exception  of  the 
carbonate  and  phosphate,  are  readily  soluble.  They  color  a  Bunsen  flame 
carmine-red  and  may  be  detected  with  the  spectroscope  by  a  strong  red 
line  to  the  left  of  the  sodium  line.  (See  frontispiece.) 

The  only  lithium  ion  known  is  Li,  for  the  reactions  of  which  any  sol- 
uble salt  will  serve. 

The  carbonate  ion,  C03,  precipitates,  from  concentrated  solutions 
only,  white  Li2C03  which  is  soluble  in  about  125  parts  of  H20. 

The  phosphate  ion,  P04,  precipitates  from  fairly  concentrated  solu- 
tions Li3PO4  which  is  soluble  in  about  2500  parts  of  H20.  It  readily 
forms  supersaturated  solutions  from  which  precipitation  is  slow.  That 
normal  lithium  phosphate  may  be  precipitated  from  solution  while  the 
normal  phosphates  of  the  other  alkali  metals  are  hydrolyzed,  is  due  to 
the  fact  that  lithium  is  somewhat  less  basic  than  the  other  members  of 
this  group  and  its  phosphates  are  much  less  soluble. 


SODIUM.  67 

SODIUM  (Natrium).    Na  —  23. 

Sodium  is  a  light,  silver-white,  soft  metal.  It  is  acted  upon  rapidly 
by  the  atmosphere,  due  chiefly  to  the  water-vapor  present,  and  requires 
to  be  kept  under  petroleum  or  some  other  liquid  which  does  not  attack 
it  and  will  protect  it  from  the  air. 

Heated  in  the  air  sodium  burns  with  an  intensely  yellow  flame,  not  to 
the  oxid,  Na2O,  but  to  the  peroxid,  Na202,  a  heavy  straw-colored  powder. 
When  thrown  into  water  it  reacts  violently  according  to  the  equation: 

Na202  +  2HOH  -»  2NaOH  +  H2O2 

Sodium  peroxid  in  water  is  equivalent,  therefore,  to  hydrogen  per- 
oxid  in  an  alkaline  solution. 

Sodium  acts  as  one  of  the  strongest  base-forming  elements.  It  is 
violently  attacked  by  all  acids,  the  concentration  of  H  ions  due  to  pure 
water  being  sufficient  to  cause  a  very  vigorous  action  at  ordinary  tem- 
peratures. Hydrogen  gas  is  liberated  and  sodium  hydroxid  formed, 
according  to  the  equation  : 


2Na  +  2HOH  -»  2NaOH 
or  writing  an  ionic  equation  : 

2Na  +  2H  +  2  OH  ->  2Na  +  2  OH  +  H2 

This  reaction  shows  that  water  acts  in  this  case  as  a  typical  acid. 
(See  page  29.)  It  may  be  regarded,  therefore,  as  the  weakest  of  all 
acids,  the  metallic  hydroxids  being  its  salts. 

So  much  energy  is  liberated  when  the  metal  passes  into  the  ionic 
condition  that  its  compounds  are  among  the  most  stable  salts  known. 
They  can  be  reduced  to  the  metal  in  the  "dry  way"  only  and  at  an  ex- 
pense of  considerable  energy. 

Sodium  salts  are  generally  colorless  and,  with  the  exception  of  a  few 
of  infrequent  occurrence,  are  readily  soluble. 

Sodium  salts  color  a  Bunsen  flame  an  intense  yellow.  In  the  spec- 
troscope, sodium  is  indicated  by  a  bright  yellow  line.  (See  frontispiece.) 
This  reaction  is  so  delicate  that  one-millionth  of  a  milligram  may  be 
readily  detected.  Hence  the  sodium  line  is  omnipresent  and  for  this 
reason  is  used  in  locating  the  position  of  other  lines  in  the  spectrum. 

The  only  known  sodium  ion  is  Na,  for  the  reaction  of  which  any  salt 
will  serve. 


68  QUALITATIVE    ANALYSIS. 

The  hydroxyl  ion,  OH,  added  to  a  solution  containing  the  Na  ion, 
produces  no  precipitate,  for  NaOH  is  soluble  in  half  its  own  weight  of 
water.  In  dilute  solutions,  sodium  hydroxid  is  almost  completely  dis- 
sociated and  is  one  of  the  strongest  bases  known.  For  this  reason  it  is 
much  used  to  introduce  OH  ions  into  solutions  containing  the  ions  of 
other  metals.  The  hydroxids  of  the  metals  generally  are  much  less 
soluble  than  NaOH  and,  with  the  exception  of  the  other  alkali  metals, 
the  hydroxid  (or  oxid)  of  each  may  be  precipitated  by  the  addition  of  a 
solution  of  sodium  hydroxid  of  suitable  concentration.  An  excess  of  the 
reagent  may  cause  the  precipitate  to  dissolve,  owing  to  the  formation  of 
soluble  compounds.  (See  lead,  antimony,  arsenic,  tin,  aluminum,  chro- 
mium, cobalt  and  zinc.) 

This  is  due  to  the  fact  that  the  hydroxids  of  certain  weak  bases,  which 
normally  separate  the  hydroxyl  ion  to  a  slight  extent,  may  also  dissociate 
into  H  ions  as  well  and  act  as  weak  acids.  The  high  concentration  of  OH 
ions  from  a  solution  of  NaOH  or  KOH  represses  the  ionization  of  these 
hydroxids  into  OH  ions  and  calls  forth  the  separation  of  H  ions  instead. 
The  strong  base  unites  with  the  weaker  hydroxid,  now  acting  as  an  acid, 
to  form  a  soluble  salt.  (See  page  22.) 

On  standing  in  the  air,  sodium  hydroxid,  like  all  strong  bases,  absorbs 
carbon  dioxid  and  passes  into  the  carbonate. 

2NaOH      C0  ->  NaC0        H0 


2  -        23          2 


For  this  reason  NaOH  always  contains  some  Na2C03. 

As  "soda  lye"  or  "caustic  soda,"  sodium  hydroxid  finds  many  ap- 
plications in  the  industries. 

The  carbonate  ion,  C03,  produces  no  precipitate,  for  Na2C03  is 
readily  soluble.  With  the  exception  of  the  other  alkalies,  the  carbonates 
of  the  metals  are  much  less  soluble  than  Na2CO3  and,  on  the  addition 
of  this  reagent,  they  (or  their  decomposition  products)  may  be  pre- 
cipitated from  solutions  of  their  soluble  salts  (arsenic  excepted,  see 
page  50  )  .  Furthermore,  since  the  sodium  salts  of  all  ordinary  acids  are 
soluble,  by  boiling  any  precipitated  substance  with  Na2C03,  the  metallic 
elements  may  be  obtained  as  carbonates  or  their  decomposition  products 
in  a  new  precipitate  while  the  acid  radicles  are  separated  as  soluble 
sodium  salts  in  the  filtrate.  Substances  so  treated  are  said  to  be 
transposed. 

In  solution,  sodium  carbonate,  like  all  other  soluble  carbonates,  is 
hydrolyzed  appreciably  (see  page  29),  and  gives  a  basic  reaction  to  lit- 


SODIUM    SALTS.  69 

mus.  Since  its  solution  contains  an  appreciable  concentration  of  OH 
ions,  it  may  cause  the  precipitation  of  the  hydroxid  as  well  as  the  car- 
bonate of  any  metal,  provided  these  two  compounds  are  soluble  to  about 
the  same  degree.  Such  is  the  case  with  several  metals.  (See  lead,  cop- 
per, bismuth,  cobalt,  nickel,  zinc  and  magnesium.)  On  treating  a  solu- 
tion of  their  soluble  salts  with  any  soluble  carbonate,  a  precipitate  falls, 
apparently  consisting  of  a  mixture  of  the  hydroxids  and  carbonates  of 
these  metals.  Such  precipitates  are  usually  called  "basic  carbonates. " 
Since  the  composition  of  any  of  these  precipitates  varies  much  with  the 
conditions  under  which  it  is  brought  down,  it  is  questionable  whether 
they  are  to  be  considered  as  true  basic  salts  (see  page  21)  or  simply 
mixtures. 

From  solution,  Na2C03  separates  with  ten  molecules  of  water  of 
crystallization.  On  gentle  heating  it  becomes  anhydrous,  after  which  it 
may  be  fused  without  decomposition,  in  this  respect  differing  from  all 
other  carbonates  except  those  of  the  other  fixed  alkalies.  (See  page  54.) 

When  fused  with  5-10  times  its  weight  of  Na2C03,  nearly  every  insol- 
uble salt  of  other  metals  is  transposed.  A  metathesis  is  started  according 
to  the  following  typical  reaction : 

Na2C03  +  PbS04  ->  Na2S04  +  PbC03  ->  PbO  +  C02 

This  reaction  is  prevented  from  coming  to  equilibrium  as  long  as  an 
appreciable  amount  of  PbS04  is  present,  on  the  one  hand,  by  the  large 
mass  of  Na2C03  present,  and,  on  the  other  hand,  by  the  fact  that  lead 
carbonate*  is  decomposed  as  fast  as  formed.  But  little  is  actually  known 
of  the  principles  underlying  such  reactions,  however,  and  although  the 
laws  governing  reactions  in  solution  seem  to  hold  under  these  conditions 
also,  they  must  be  applied  with  caution  because  of  the  lack  of  definite 
data. 

In  making  fusions,  equal  parts  of  Na2C03  and  K2C03  are  frequently 
used  instead  of  Na2CO3  alone,  inasmuch  as  the  mixture  fuses  at  a  lower 
temperature  than  either  pure  salt.  This  phenomenon  follows  the  general 
rule  that  the  melting  point  of  any  substance  is  lowered  by  the  admixture 
of  any  substance  with  it. 

Under  the  names  "soda,"  "sal  soda"  and  "washing  soda,"  Na2C03 
finds  many  applications. 

*  When  insoluble  salts  of  barium,  strontium  or  calcium  are  fused  in  this  man- 
ner, the  reaction  goes  in  the  same  direction  forming  carbonates,  but  the  temperature 
of  an  ordinary  fusion  is  not  sufficient  to  decompose  them  into  CO2  and  the  oxids  (see 
pages  99,  98,  and  95). 


70  QUALITATIVE    ANALYSIS. 

The  nitrate  ion,  N03,  causes  no  precipitation,  for  all  normal  nitrates 
are  readily  soluble.  When  heated,  sodium  nitrate,  like  all  nitrates, 
decomposes  into  oxygen  and  oxids  of  nitrogen  and  of  the  metal  (see  page 
150).  For  this  reason  it  is  much  used  as  an  oxidizing  agent  in  reac- 
tions performed  in  the  dry  way.  Heated  gently  with  some  mild  reduc- 
ing agent,  it  breaks  down  into  sodium  nitrite,  NaN02,  from  which  all 
other  nitrites  may  be  formed.  As  ' '  Chili  saltpeter  "  or  "  caliche, ' '  NaN03, 
is  used  extensively  in  making  HN03,  KN03  and  in  fertilizers. 

The  sulfid  ion,  S,  causes  no  precipitate,  for  the  alkaline  sulfids  are 
readily  soluble. 

The  sulfate  ion,  S04,  causes  no  precipitate,  for  Na2S04  is  readily 
soluble.  From  solutions,  it  crystallizes  at  ordinary  temperatures  with 
ten  molecules  of  water  and  is  known  as  "Glauber's  salt."  It  is  much 
used  in  the  manufacture  of  "soda"  and  glass. 

The  chromate  ion,  Cf04,  causes  no  precipitate.  Since  Na2Cr04. 
10H20  is  very  deliquescent  as  well  as  very  soluble,  it  can  not  be  easily 
preserved,  hence  the  potassium  salt,  K2Cr04,  is  more  often  used  for  the 
purpose  of  introducing  CrO4  ions  into  solutions  containing  metallic  ions. 

For  similar  reasons  K2Cr207  is  preferable  to  Na2Cr2O7  •  2H2O  for 
general  laboratory  use. 

The  chlorid  ion,  01,  produces  no  precipitate,  for  the  halids  are  all 
soluble,  increasing  in  degree  as  the  atomic  weight  of  the  halogen  rises 
(see  table,  page  71).  From  hot  solutions  all  crystallize  anhydrous. 

Sodium  cyanid,  NaCN,  is  very  soluble  and  very  deliquescent. 

The  phosphate  ion,  P04,  unites  with  the  Na  ion  to  form  soluble  salts 
of  three  distinct  types.  Tertiary  sodium  phosphate,  Na3P04,  the  normal 
salt,  is  formed  only  in  the  * '  dry  way ' '  and,  when  put  into  solution,  hydro- 
lyzes  largely  into  the  secondary  salt,  Na2HP04,  and  NaOH,  for  the  solu- 
tion reacts  strongly  alkaline  and  crystallizes  out  Na2HP04.12H20.  This 
acid  disodium  phosphate,  the  secondary  salt,  is  the  substance  to  which 
reference  is  made  when  the  term,  "sodium  phosphate,"  is  used.  To 
such  a  slight  degree  does  phosphoric  acid,  H3PO4,  separate  all  three  of 
its  hydrogen  atoms  as  ions  (see  page  155)  that  the  reaction  to  litmus 
of  this  acid  salt  is  not  what  would  naturally  be  expected,  for  hydrolysis 
of  Na2HP04  into  NaH2PO4  and  NaOH  puts  more  OH  ions  into  solution 
than  the  number  of  H  ions  formed  by  dissociation.  Hence  it  has  a  slight 
basic  reaction  to  litmus. 

The  acid  pyroantimonate  ion,  H2Sb20T,  produces,  in  concentrated 
neutral  solutions,  a  white  crystalline  precipitate  of  Na2H2Sb2O7,  which 
is  soluble  in  about  400  parts  of  H2O.  None  of  the  other  alkalies  give  a 


SODIUM    SALTS. 


71 


similar  reaction.  Metals  other  than  the  alkalies  must  be  absent,  as  they 
generally  produce  amorphous  precipitates.  This  is  the  only  form  in 
which  sodium  may  be  separated  by  precipitation  in  aqueous  solution 
from  the  other  alkali  metals.  The  method  is  scarcely  to  be  recommended 
for  general  use,  as  the  solution  of  the  precipitating  reagent,  K2H2Sb207, 
must  be  prepared  fresh  for  each  test. 

Inasmuch  as  the  anions  formed  by  the  various  acids  are  frequently 
introduced  into  solutions  by  means  of  their  sodium  salts,  a  list  of  the 
salts  of  the  commoner  acids  is  appended  together  with  a  statement  of 
some  of  their  properties,  i.  e.,  color,  character,  approximate  solubility 
and  the  formula  for  the  ordinary  salt. 


Acid. 

Formula. 

Color. 

Character. 

Solubility.* 

Acetic 

Na^dH8O2).3H2O 

White 

Efflorescent 

3 

Arsenic 

Na2HAsO4.12H2O 

White 

Efflorescent 

1.75 

Arsenous   (meta) 

NaAsO2 

White 

Poor  crystals 

Very  sol. 

Boric   (tetra) 

Na2B4O7.10H2O 

White 

Efflorescent 

14 

Carbonic 

Na2CO8.10HaO 

White 

Efflorescent 

1.1 

Carbonic,  anhydrous 

Na2C03 

White 

Stable 

5 

Chloric 

NaClO3 

White 

Deliquescent 

3 

Chromic 

Na2CrO4.10H2O 

Yellow 

Deliquescent 

0.1 

Chromic   (pyro)  — 

'  '  dichromic  '  } 

Na,Cr2OT.2H2O 

Bed 

Deliquescent 

0.9 

Cyanic 

NaOCN 

White 

Stable 

Sol. 

Ferricyanic 

Na,(FeC,N8).H20 

Bed 

Deliquescent 

5 

Ferrocyanic 

Na4(FeC6N.).12H20 

Yellow 

Efflorescent 

4.5 

Hydrobromie 

NaBr 

White 

Stable 

1.15 

Hydrochloric 

NaCl 

White 

Stable 

3 

Hydrocyanic 

NaCN 

White 

Deliquescent 

Very  sol. 

Hydrofluoric 

NaF 

White 

Stable 

25 

Hydriodic 

Nal 

White 

Stable 

0.6 

Hydrosilicofluoric 

Na2SiF, 

White 

Stable 

150 

Hydrosulfuric 

Na2S 

Straw 

Deliquescent 

Sol. 

Hypochlorous 

NaOCl 

White 

Known  only 

in  solution 

Nitric 

NaN08 

White 

Deliquescent 

1.2 

Nitrous 

NaNO, 

Straw 

Stable 

1 

Oxalic 

Na.,C2O4 

White 

Stable 

30 

Perchloric 

NaClO4 

White 

Deliquescent 

Very  sol. 

Permanganic 

NaMnO4 

Purple 

Deliquescent 

Very  sol. 

Phosphoric 

Na2HPO4.12H2O 

White 

Efflorescent 

10 

Silicic   (meta) 

Na,SiO3 

White 

Stable 

Very  sol. 

Sulfuric 

Na.,SO4.10H2O 

White 

Efflorescent 

2 

Sulfurous 

Na2SO3.8H2O 

White 

Efflorescent 

4.3 

Thiocyanic 

NaSCN 

White 

Deliquescent 

Very  sol. 

Thiosulfuric 

Na2S2O3.5H,O 

White 

Efflorescent 

1.5 

Tartaric 

Na2(C4H406).2H20 

White 

Stable 

2.5 

*  The  solubility  is  expressed  in  grams  or  cubic  centimeters  of  water  necessary  to 
dissolve  one  gram  of  the  salt. 


72  QUALITATIVE    ANALYSIS. 

POTASSIUM  (Kalium).    K  — 39.2. 

Potassium  is  a  soft,  silvery  metal  with  a  bluish  tint,  in  its  properties 
and  reactions  closely  resembling  sodium.  (See  page  67.) 

Potassium  salts  are  generally  colorless  and  usually  somewhat  less 
soluble  than  the  corresponding  sodium  compounds. 

Potassium  salts  give  to  a  Bunsen  flame  a  very  characteristic  color 
which  is  usually  described  as  lavender  or  violet.  Spectroscopically  ex- 
amined, the  flame  shows  a  red  line  considerably  to  the  left  of  the  sodium 
line  and  a  violet  line  far  to  the  right  at  the  other  end  of  the  spectrum 
(see  frontispiece).  These  lines  are  often  detected  with  some  difficulty 

by  the  untrained  observer. 

+ 

The  only  known  potassium  ion  is  K,  for  the  reactions  of  which  any  of 
the  more  soluble  salts  will  serve.  The  compounds  of  potassium  so  closely 
resemble  those  of  sodium  that  the  more  important  differences  only  will 
be  considered. 

Potassium  hydroxid,  KOH,  is  possibly  a  trifle  more  strongly  basic 
than  sodium  hydroxid.  For  most  analytical  reactions  these  reagents 
may  be  used  interchangeably.  As  "potash  lye"  or  "caustic  potash," 
it  finds  less  extensive  application  in  the  industries  than  does  "caustic 
soda." 

Potassium  carbonate,  K2C03,  or  "potash,"  is  deliquescent,  absorbing 
moisture  readily  from  the  air,  unlike  Na2C03. 

Potassium  nitrate,  KN03,  "saltpeter,"  unlike  the  sodium  salt,  is  not 
deliquescent  and,  therefore,  is  used  in  the  manufacture  of  explosives. 

Potassium  cyanid,  KCN,  is  more  commonly  used  in  the  extraction  of 
gold  and  as  a  bath  in  electroplating  with  the  precious  metals  than  is  the 
sodium  salt. 

Potassium  sulfate,  K2S04,  crystallizes  anhydrous  from  its  solutions. 
The  acid  salt,  KHS04,  melts  at  200°,  losing  water  and  changing  into 
the  pyrosulfate,  K2S207.  On  further  heating  this  compound  breaks  up 
into  the  normal  sulfate,  K2S04,  and  S03.  Since  nascent  sulfur  trioxid, 
thus  formed,  acts  vigorously  upon  many  compounds,  fusion  with  acid 
potassium  sulfate  is  used  to  decompose  silicates  and  other  compounds 
not  readily  attacked  by  acids. 

Potassium  chromate,  K2Cr04,  as  well  as  the  dichromate,  K2Cr207, 
crystallize  anhydrous  from  their  solutions.  They  are  stable  under  ordi- 
nary conditions  and,  therefore,  readily  preserved.  They  are  the  reagents 
ordinarily  used  for  introducing  CrO4  and  Cr207  ions  into  solution. 

Potassium  chlorate,  KC103,  breaks  down  readily  on  heating  into  KC1 
and  oxygen,  hence  is  much  used  as  an  oxidizing  agent,  finding  especial 


POTASSIUM.  73 

application  in  matches  and  pyrotechnics.  On  gentle  heating,  it  loses 
part  only  of  its  oxygen,  forming,  at  first,  potassium  perchlorate,  KC1O4, 
according  to  the  reaction: 

2KC103  ->  KC104  +  02  +  KC1 

Since  KC104  is  soluble  in  about  50  parts  of  H20,  whereas  NaC104  is 
much  more  soluble,  a  partial  precipitation  of  KC104  may  be  obtained  by 
treating  a  concentrated  solution  containing  the  K  ion  with  a  saturated 
solution  of  NaC104.  The  precipitation  is  made  more  complete  by  the 
addition  of  alcohol. 

Potassium  tartrate,  K2(C4H406),  is  readily  soluble.  Acid  potassium 
tartrate,  KH(C4H406),  requires  about  90  parts  of  H20  to  dissolve  it. 
In  50  per  cent,  alcohol,  however,  it  is  quite  insoluble.  The  precipitation 
must  be  made  in  a  solution  slightly  acid  with  acetic  or  some  other  weak 
acid,  for  the  salt  is  soluble  in  the  strong  acids  and  in  alkalies,  in  the 
latter  case  forming  the  readily  soluble  normal  salt.  Other  than  alkali 
metals  must  be  absent,  as  they  form  insoluble  precipitates  under  the 
same  conditions.  The  ammonium  ion,  NH4,  if  present,  will  give  a  sim- 
ilar precipitate  but,  on  ignition,  will  be  driven  off  entirely  whereas  the 
potassium  salt  will  be  changed  to  the  carbonate  as  is  the  rule  with  salts 
of  organic  acids  generally. 

2KH(C4H406)  +  502  ->  K2C03  +  5#20  +  7C02 

"Rochelle  salt,"  KNa(C4H406).4H20,  is  soluble  in  15  parts  of  H20. 

Potassium  platinic-chlorid,  K2PtClc,  soluble  in  about  100  parts  of 
H20,  is  precipitated  as  a  yellow  crystalline  precipitate  from  concen- 
trated solutions  containing  the  K  ion  by  the  addition  of  H2PtCl6.  In 
alcohol  it  is  much  less  soluble.  The  customary  method  of  procedure  is 
to  evaporate  almost  to  dryness  the  solution  containing  the  K  and  PtCl6 
ions.  Extraction  with  alcohol  dissolves  the  soluble  portion  of  the  resi- 
due, leaving  behind  the  potassium  salt.  The  ammonium  ion,  NH4,  if 
present,  gives  a  similar  precipitate  but,  on  ignition,  is  decomposed  and 
the  ammonium  radicle  completely  driven  off.  Potassium  platinic- 
chlorid,  on  ignition,  is  decomposed  and  leaves  potassium  chlorid  and  a 
residue  of  platinum  according  to  the  reaction : 

K2PtCl6  ->  2KC1  +  Pt  +  2C12 

Potassium  cobalti-nitrite,  K3Co(NO2)6,  is  precipitated  from  very 
concentrated  solutions  containing  the  K  ion  and  acidified  with  acetic 
acid,  by  the  addition  of  a  solution  of  a  cobaltous  salt  and  solid  sodium 


74 


QUALITATIVE    ANALYSIS. 


nitrite.     Nitrous  acid,  HN02,  set  free  by  the  acetic  acid,  first  oxidizes 

•H-  +++ 

the  cobaltous  ion,  Co,  to  the  cobaltic  ion,  Co,  which  then  unites  with  the 
K  ions  and  N02  ions  to  form  the  yellow  K3Co(N02)6  (see  page  -  — ). 
The  ammonium  ion,  NH4,  must  be  absent,  as  it  forms  a  similar 
precipitate. 

Potassium  silicofluorid,  K2SiF6,  is  a  white  crystalline  salt,  precipi- 
tated from  slightly  acid  solutions  containing  the  K  ion  on  the  addition 
of  the  SiF6  ion.  It  is  much  less  soluble  than  the  sodium  salt,  requiring 
about  825  parts  of  H20  to  dissolve.  An  equal  volume  of  alcohol  precipi- 
tates it  from  aqueous  solution  almost  completely. 

As  potassium  salts  are  frequently  used  to  introduce  anions  into  other 
solutions,  a  list  of  the  ordinary  salts  of  the  commoner  acids  is  appended, 
together  with  the  chemical  formula,  color,  character  and  approximate 
solubility  of  each. 


Acid. 

Formula. 

Color. 

Character. 

Solubility* 

Acetic 

K(C2H302) 

White 

Poor  crystals 

0.5 

Arsenic 

K,HAsO4 

White 

Stable 

Sol. 

Arsenous  (meta) 

KAsO, 

White 

Stable 

Sol. 

Boric  (meta) 

KBO2 

WThite 

Stable 

Sol. 

Carbonic 

K2CO3 

WThite 

Deliquescent 

0.9 

Chloric 

KC1O3 

White 

Stable 

16 

Chromic 

KoCrO4 

Yellow 

Stable 

1.75 

Chromic   (pyro)  —  ' 

K2Cr2OT 

Bed 

Stable 

10 

'*  dichromic" 

Cyanic 

KOCN 

White 

Stable 

Sol. 

Ferricyanic 

K3(FeC6N6) 

Red 

Stable 

2.5 

Ferrocyanic 

K4(FeC6N6).3H20 

Yellow 

Stable 

3.5 

Hydrobromic 

KBr 

White 

Stable 

1.5 

Hydrochloric 

KC1 

White 

Stable 

3 

Hydrocyanic 

KCN 

White 

Deliquescent 

Very  sol. 

Hydrofluoric 

KF 

White 

Deliquescent 

Very  sol. 

Hydriodic 

KI 

White 

Stable 

0.7 

Hydrosilicofluorie 

K2SiF8 

White 

Stable 

825 

Hydrosulfuric 

K2S 

Straw 

Deliquescent 

Sol. 

Hypochlorous 

KC10 

White 

Known  only 

in  solution 

Nitric 

KN03 

White 

Stable 

3.3 

Nitrous 

KN02 

Straw 

Deliquescent 

Sol. 

Oxalic 

K,C204.H2O 

White 

Stable 

3 

Perchloric 

KC104 

White 

Stable 

60 

Permanganic 

KMnO4 

Purple 

Stable 

15 

Phosphoric 
Silicic   (meta) 

K2HP04 
K2Si03 

White 
White 

Deliquescent 
Stable 

Very  sol. 
Very  sol. 

Sulfuric 

K2SO4 

White 

Stable 

10 

Sulfurous 

K,SO3.2H2O 

White 

Deliquescent 

1 

Thiocyanic 

KCNS 

White 

Deliquescent 

0.5 

Thiosulfuric 

K2S203 

White 

Deliquescent 

Very  sol. 

Tartaric 

K2(C4H406).H20 

White 

Stable 

0.6 

*  The  solubility  is  expressed  in  grams  or  cubic  centimeters  of  water  necessary  to 
dissolve  one  gram  of  the  salt. 


AMMONIUM.  75 

AMMONIUM,  NH4. 

Ammonium  is  a  combination  of  nitrogen  and  hydrogen  in  the  pro- 
portions of  one  volume  of  the  former  to  four  volumes  of  the  latter.  It 
exists  in  a  great  number  of  salts  and  other  compounds  in  which  it  takes 
the  place  of  a  metal  and  shows  the  properties  of  a  base-forming  element. 
It  is  equally  well  known  as  a  univalent  kathion,  separating  at  the  nega- 
tive pole  in  electrolysis  as  do  the  metals.  It  is  not  known  free,  however, 
for  as  soon  as  it  is  separated  from  any  of  its  compounds,  it  decomposes 
into  ammonia,  NH3,  and  hydrogen. 

Ammonium  salts  resemble  those  of  potassium  very  closely  in  their 
physical  properties.  Unlike  the  latter,  they  are  all  volatile,  volatilization 
being  accompanied  by  more  or  less  decomposition  which  nray  be  either 
temporary  or  permanent. 

Salts  of  acids  containing  no  oxygen  are  volatile  (generally  subliming) 
with  dissociation  into  ammonia  and  the  acid,  which  recombine  on  cooling. 


Salts  of  acids  containing  oxygen  which  are  not  oxidizing  agents  split 
off  the  ammonia,  gradually  forming  in  the  case  of  polybasic  acids,  acid 
salts  first,  which  later  decompose  into  ammonia  and,  generally,  water  and 
the  acid  anhydrid:  * 

(NH4)2S04  -»  NHS  +  H(NH4)S04 

H(NH4)S04  -»  NH3  +  H20  +  S03 

Salts  of  acids  which  are  oxidizing  agents  decompose  by  an  internal 
oxidation,  the  oxidizing  and  reducing  agents  both  being  contained  in  the 
compound  itself.  "When  ammonium  nitrite  is  heated,  the  nitrite  radicle 
oxidizes  the  ammonium  radicle  as  indicated  in  the  equation  : 


Since  all  ammonium  salts  may  be  removed,  if  subsequently  necessary, 
by  moderate  heating,  they  are  much  used  for  introducing  various  anions 
into  solutions  containing  metallic  ions. 

Ammonium  salts  give  no  color  to  a  Bunsen  flame. 

The  only  known  ammonium  ion  is  NH4. 

Ammonium  hydroxid,  NH4OH,  is  known  only  in  solution,  for,  on 
attempting  to  isolate  it,  decomposition  into  NH3  and  H20  takes  place. 
For  this  reason  it  is  known  as  the  ''volatile  alkali."  The  reaction  be- 


76  QUALITATIVE    ANALYSIS. 

tween  ammonia  and  water  (see  page  147)  is  a  typical  equilibrium  re 
action,  as  is  shown  by  the  equation  : 


In  solution,  there  is  always  a  large  portion  of  the  gas  dissolved  in 
the  water  that  is  not  united  with  it  in  the  form  of  NH4OH,  hence  the 
solution  is  frequently  known  as  "ammonia  water"  as  well  as  ammonium 
hydroxid.  This  free  ammonia  shows  a  special  tendency  to  unite  with 
various  simple  metallic  ions  to  form  complex  ions. 

Ammonium  hydroxid  separates  the  OH  ion  to  a  considerably  greater 
degree  than  do  most  metallic  hydroxids.  Compared  with  the  alkaline 
hydroxids  or  even  the  hydroxids  of  the  alkaline  earth  metals,  it  is  only 
a  weak  base,  however.  A  solution  of  potassium  or  sodium  hydroxid  of 
ordinary  concentration  contains  more  than  sixty  times  as  many  OH  ions 
as  a  solution  of  ammonium  hydroxid  of  the  same  concentration.  Al- 
though these  solutions  are  frequently  used  interchangeably  for  the  intro- 
duction of  OH  ions,  it  is  only  to  be  expected  that  the  reactions  would 
not  be  exactly  the  same  in  every  case.  (See  different  action  of  NH4OH 
from  KOH  or  NaOH  on  lead,  tin  and  zinc  hydroxids.) 

Because  sodium  or  potassium  hydroxid  dissociates  so  much  more 
than  ammonium  hydroxid,  salts  of  the  latter  are  decomposed  by  the 
former  reagents.  Ammonium  chlorid,  as  an  example,  on  going  into  solu- 

4* 

tion,  dissociates  almost  completely,  giving  high  concentrations  of  NH4 
and  of  Cl  ions.  On  the  addition  of  KOH,  high  concentrations  of  K  and 
of  OH  ions  are  introduced.  Of  the  compounds  which  may  be  formed  by 
these  four  different  ions,  all  dissociate  readily  except  NH4OH.  As  this  is 

formed,  almost  the  whole  of  it  remains  undissociated  and,  therefore,  cuts 

+ 
down  the  concentration  of  NH4  and  of  OH  ions  and  removes  them  from 

the  sphere  of  action.  On  warming  such  a  solution,  NH4OH  is  decomposed 
and  NH3  passes  off,  leaving  in  solution  only  K  and  Cl  ions  to  form  KC1. 
This  reaction  is  ordinarily  used  as  a  test  for  ammonium  salts. 

Ammonium  carbonate,  (NH4)2C03,  is  a  white  substance,  soluble  in 
about  one  part  of  water,  used  frequently  as  a  reagent  for  the  introduction 
of  CO3  ions  into  solutions  containing  metallic  ions.  Like  all  soluble  car- 
bonates, it  hydrolyzes  greatly  into  NH4OH  and  H2CO3,  both  of  which 
decompose  into  gaseous  products.  Hence  a  solution  of  ammonium  car- 
bonate contains  NH4OH  and  NH3.  Hydrolysis  increases  as  the  tempera- 
ture is  raised,  hence  by*boiling,  (NH4)2C03  may  be  entirely  decomposed 
and  removed  from  solution. 


AMMONIUM    SALTS.  77 

Ammonium  carbonate  splits  off  part  of  its  ammonia  even  at  ordinary 
temperatures  and  forms  the  acid  salt.  For  this  reason,  the  commercial 
product  usually  consists  largely  of  acid  ammonium  carbonate, 
H(NH4)C03.  Ammonium  carbonate  also  loses  water  readily  forming 
ammonium  carbamate  according  to  the  reaction: 

H4N  —  0\  H4N  — 0\ 

>C  =  0->  >C  =  0  +  H20 

H4N  —  O/  H2N 

This  product  is  also  found  in  the  commercial  article.  By  adding 
NH4OH  to  a  solution  of  commercial  ammonium  carbonate,  however,  no 
new  substance  is  introduced  and  the  acid  salt  is  changed  into  the  normal 
salt.  Under  the  same  conditions,  also,  the  carbamate  reverts  to  the 
carbonate. 

Ammonium  sulfid,  (NH4)2S,  may  be  formed  as  a  white  solid  by  bring- 
ing together  NH3  and  H2S  gases  at  low  temperatures.  At  ordinary  tem- 
peratures, it  is  decomposed  almost  entirely  into  the  constituents.  In 
solution  it  is  fairly  stable  when  protected  from  the  action  of  the  atmo- 
sphere and  of  light.  It  is  usually  formed  by  saturating  a  volume  of 
strong  NH4OH  with  H2S  which  forms  acid  ammonium  sulfid  according 
to  the  reaction : 

NH4  —  OH  +  H  —  S  —  H->NH4  —  S  —  H  +  H20 

By  the  addition  of  another  volume  of  NH4OH  the  normal  salt  is  formed. 

A  fresh  solution  of  (NH4)2S  is  colorless  and  reacts  alkaline,  owing 
to  hydrolysis  into  NH4OH  and  H2S.  On  standing  in  the  light  and  air,  it 
becomes  yellow.  The  H2S  oxidizes  slowly  into  H20  and  S,  the  latter 
substance  dissolving  in  and  reacting  with  (NH4)2S  to  form  some  of  the 
ammonium  polysulfids,  of  which  several  are  known,  such  as  (NH4)2S2, 
(NH4)2S4  and  so  forth.  A  solution  of  these  polysulfids  is  known  as 
"yellow  ammonium  sulfid." 

Ammonium  sulfid  is  much  used  for  the  introduction  of  S  ions  into 
solution  because  it  furnishes  a  concentration  of  S  ions  much  greater  than 
can  be  obtained  by  a  saturated  solution  of  H2S.  Moreover,  it  may  be 
removed  from  solution  simply  by  boiling,  which  is  not  true  of  the 
sulfids  of  the  fixed  alkalies. 

Ammonium  polysulfid  is  used  to  dissolve  certain  metallic  sulfids 
(see  antimony,  arsenic,  tin,  gold  and  platinum)  with  which  it  unites  to 
form  soluble  salts  of  thioacids. 


78  QUALITATIVE    ANALYSIS. 

Ammonium  nitrate,  NH4N03,  is  a  colorless,  crystalline,  deliquescent 
salt,  soluble  in  about  half  its  weight  of  water.  On  heating  it  decom- 
poses, the  nitric  acid  radicle  oxidizing  the  ammonium  radicle  according 

to  the  equation : 

NH4N03  -»  N20  +  2H20 

In  the  presence  of  oxidizable  material,  the  reaction  does  not  form 
N20  but  N2,  the  oxygen  being  taken  up  by  the  reducing  agent.  For  this 
reason,  it  is  often  used  in  fusions  to  remove  organic  matter. 

Ammonium  acetate;  NH4(C2H302),  is  a  colorless,  crystalline  deli- 
quescent substance  which  is  very  soluble  in  water. 

Ammonium  chlorid,  NH4C1,  is  a  colorless,  crystalline  substance, 
soluble  in  2.5  parts  of  H20.  It  unites  with  the  salts  of  many  metals 
to  form  double  salts  and  is  thus  used  to  prevent  the  precipitation 
of  metallic  hydroxids  by  NH4OH.  ( See  magnesium,  ferrous  iron,  cobalt, 
nickel,  manganese  and  zinc.  See  also  page  32.) 

The  ammonium  phosphates  are  unimportant  and  closely  resemble  the 
salts  of  sodium  and  potassium.  Microcosmic  salt,  HNa(NH4)P04.4H20, 
is  interesting  because  of  its  decomposition  on  heating  into  sodium  meta- 
phosphate,  ammonia  and  water. 

H  — 0  0 

\  / 

H4N  —  0— P  =  O^Na  —  O  —  P       +NH3  +  H20 

/  \ 

Na  — 0  0 

Sodium  meta-phosphate  is  used  frequently  in  fusions.  Being  a  flux 
for  all  basic  oxids,  it  leaves  non-volatile,  infusible,  acid  anhydrids,  such 
as  silica,  undissolved  and  floating  about  in  the  clear  melt.  (See  test  for 
silicic  acid,  page  260.) 

Ammonium  oxalate,  (NH4)2C204-  H20,  is  a  colorless  crystalline  body, 
soluble  in  about  23  parts  of  H2O.  It  is  generally  used  for  the  introduc- 
tion of  the  oxalate  ion  into  solutions. 

A  few  less  soluble  complex  salts  of  ammonium  are  known,  such  as  the 
acid  tartrate,  platinic-chlorid  and  cobalti-nitrite.  They  resemble  the 
potassium  salts  very  closely.  ( See  page  74. ) 

Ammonium  silicofluorid,  (NH4)2SiF6,  is  soluble  in  about  5  parts 
of  H20. 

Ammonium  molybdate,  (NH4)2Mo04,  is  a  white  crystalline  substance 
without  water  of  crystallization.  It  is  soluble  in  about  5  parts  of  ELO 
and  is  used  as  a  reagent  for  detecting  phosphoric  and  arsenic  acids. 


THE    COPPER    GROUP.  79 

Group  1,  Second  Type. 
Copper  63.6;        Silver  107.9;        Gold  197.2. 

The  members  of  this  group  are  heavy,  malleable  metals  which  are 
good  conductors  of  heat  and  electricity.  They  are  found  in  nature  both 
free  and  in  combination  with  other  elements  and  they  are  generally 
stable  under  ordinary  conditions.  They  show  but  little  tendency  to  pass 
into  the  ionic  condition,  hence  are  readily  reduced  from  their  compounds. 

The  group,  as  a  whole,  is  not  well  characterized,  however.  The  sim- 
ilarity of  these  elements  to  each  other  is  not  striking  nor  is  the  resem- 
blance of  this  half  to  the  other  half  of  the  same  group  sufficiently  obvious. 
Copper  and  gold  resemble  elements  in  other  groups  quite  closely  and 
their  right  to  a  position  in  this  group  is  not  beyond  question.  Silver  is 
a  typical  metal,  acting  as  a  fairly  strong  base-forming  element.  Its 
compounds  are  colorless  and  resemble  those  of  the  alkali  metals  and,  like 
them,  silver  is  consistently  univalent. 

COPPER  ( Cuprum) .     Cu  —  63.6. 

Copper  is  a  rather  soft,  very  malleable  metal,  possessing  a  character- 
istic red  color.  It  tarnishes  in  moist  air,  forming  a  green  basic  carbon- 
ate sometimes  known  as  verdigris.  At  higher  temperatures,  it  unites 
readily  with  oxygen,  giving  rise  to  the  black  oxid,  CuO. 

The  metal  is  attacked  but  little  by  HC1  or  by  dilute  H2S04.  This 
fact  is  readily  explained  by  reference  to  the  Potential  Series  (given  on 
page  327),  which  shows  that  hydrogen  passes  into  the  ionic  condition 
more  readily  than  copper  does.  Consequently,  when  metallic  copper  is 
brought  into  contact  with  hydrogen  ions,  there  is  little  tendency  for  the 
latter  to  give  up  their  electric  charges  to  the  former  and,  therefore,  the 
metal  does  not  dissolve  appreciably. 

Oxidizing  agents,  however,  convert  the  metal  into  the  oxid  and  this 
will  dissolve  readily  in  acids.  In  this  way  the  metal  may  be  obtained 
in  solution.  Nitric  acid  acts  both  as  an  oxidizing  agent  and  as  an  acid 
and  for  this  reason  readily  dissolves  those  metals  which,  like  copper, 
show  less  tendency  to  pass  into  the  ionic  condition  than  hydrogen  does. 
Since  hot,  concentrated  H2SO4  may  act  as  an  oxidizing  agent  (see  page 
187),  it,  also,  will  dissolve  copper  and  similar  metals. 

By  further  reference  to  the  Potential  Series,  it  will  be  seen  that  most 
metals  tend  to  pass*  into  the  ionic  condition  more  readily  than  copper 
does.  They  are,  therefore,  said  to  be  more  positive  metals.  If  one  of 
these  metals  is  introduced  into  a  solution  of  a  salt  of  copper,  the  charge 
passes  from  the  copper  ions  to  the  metal.  Copper  is  precipitated  from 


80  QUALITATIVE    ANALYSIS. 

solution  in  metallic  form  while  the  more  positive  metal  passes  into  solu- 
tion in  the  form  of  ions,  as  the  equations  indicate. 

CuS04  +  Zn  ->  ZnS04  +  Cu  or,  simplified,  Cu  +  Zn  ->  Cu  +  Zn 

Aluminum,  iron  and  other  metals  will  precipitate  metallic  copper 
from  solutions  of  its  salts  because  of  the  decidedly  greater  tendency  of 
these  metals  to  become  ions. 

All  elements  which,  like  copper,  show  little  tendency  to  pass  into  the 
ionic  condition,  are  readily  reduced  from  their  salts  to  the  metal  and, 
conversely,  those  elements  which  pass  readily  into  the  ionic  condition  are 
reduced  only  with  difficulty. 

When  heated  before  the  blowpipe  with  Na2C03  on  charcoal,  all  cop- 
per salts  are  reduced,  but  the  metal  shows  little  tendency  to  come  to- 
gether in  a  single  globule.  Under  favorable  conditions  a  minute  red 
button  may  be  formed  without  a  coating  on  the  charcoal.  It  is  readily 
soluble  in  warm  HNO3,  giving  a  blue  solution,  especially  after  the  addi- 
tion of  NH4OH. 

Copper  salts  generally  impart  a  green  color  to  a  Bunsen  flame,  very 
similar  to  the  color  produced  by  barium  salts  or  boric  acid.  Cupric 
chlorid,  however,  gives  a  beautiful  blue  color  which  can  hardly  be  mis- 
taken. By  moistening  the  substance  to  be  tested  with  HC1  and  heating 
in  a  loop  of  platinum  wire,  the  blue  flame  will  usually  be  seen.  A  better 
way  is  to  pass  HC1*  gas  into  one  of  the  air-holes  at  the  base  of  a  Bunsen 
burner  (which  must  not  be  made  of  brass)  and  heat  the  substance  on  a 
platinum  wire  in  the  flame,  when,  if  copper  be  present,  the  blue  flame 
will  be  observed. 

Copper  acts  as  a  weak  base-forming  element.  Its  salts  hydrolyze 
considerably  and,  in  solution,  exhibit  a  tendency  to  deposit  insoluble 

basic  salts.     Two  kinds  of  simple  copper  ions  are  known,  the  cupric  ion, 
++  + 

Cu,  and  the  cuprous  ion,  (Cu)2.  These  resemble  considerably  the  cor- 
responding mercury  ions  (see  page  107).  There  are  also  several  com- 
plex ions,  formed  especially  with  NH3  and  the  CN  ion,  which  will  be 
considered  in  their  proper  places. 

From  the  simple  ions,  two  series  of  salts  are  formed  which  differ 
decidedly  in  their  properties.  The  more  stable  cupric  salts  are  colored 
usually  blue  or  green,  the  nitrate,  sulfate,  acetate  and  chlorid  being  solu- 
ble. No  cuprous  salts  of  acids  containing  oxygen  are  known.  The 

*  A  little  HC1  generator  suitable  for  this  purpose  may  be  made  by  heating  con- 
centrated HC1  in  a  test-tube  provided  with  a  delivery-tube  leading  into  one  of  the  air- 
holes at  the  base  of  the  Bunsen  burner. 


CUPKIC    SALTS.  81 

halids  are  colorless  and  highly  insoluble  in  water.  They  tend  to  take  up 
oxygen  and  pass  into  the  higher  form.  The  reactions  of  the  two  ions 
will  be  considered  separately. 

The  Cupric  Ion,  Cu. 

Cupric  nitrate  (or  sulfate  or  chlorid)  is  a  good  source  for  Cu  ions. 
It  is  deep  blue  in  color  and  crystallizes  from  solution  with  3  or  6  H20. 
Both  these  substances  are  deliquescent  and  therefore  can  not  be  pre- 
served as  satisfactorily  as  salts  which  do  not  take  up  water  vapor  from 
the  air.  Both  are  extremely  soluble.  At  38°,  Cu(N03)2.6H20  will  dis- 
solve in  its  own  water  of  crystallization. 

The  hydroxyl  ion,  OH,  precipitates  floccy,  blue  Cu(OH)2,  which 
tends  to  lose  water  slowly  at  ordinary  temperatures  and,  on  boiling, 
changes  to  black  CuO. 

In  the  presence  of  NH3,  the  precipitate  first  formed  is  redissolved, 
owing  to  the  formation  of  a  complex  cupric-ammonia  ion,  Cu(NH3)4, 
the  compounds  of  which  are  soluble.  This  ion  is  somewhat  unstable  and 
the  concentration  of  Cu  ions  from  its  decomposition  is  sufficient  to  cause 
a  partial  precipitation  of  CuS  when  the  S  ion  is  introduced.  On  the 
addition  of  acids  the  deep  blue  color  characteristic  of  the  Cu(NH3)4 
ion  disappears  because  of  the  decomposition  of  the  complex  ion  caused  by 
the  union  of  the  H  ion  with  the  NH3  to  form  the  ammonium  ion,  NH4. 
The  addition  of  KCN  bleaches  the  solution  also,  owing  to  the  formation  of 
colorless  compounds  as  the  following  typical  reactions  indicate : 

Cu(NH3)4S04  +  2KCN  +  4H20 -»  CuC2N2  +  K2S04  +  4NH4OH 

2CuC2N2-»(CuCN)2  +  (CN)2 
(CuCN)2  +  6KCN-*  2K8(CuC4N4) 

Cupric  hydroxid  dissolves  in  solutions  of  alkaline  cyanids,  forming 
complex  cyanid  ions. 

In  the  presence  of  tartaric  acid,  citric  acid,  sugar  or  other  organic 
substances  containing  hydroxyl  radicles,  no  precipitation  takes  place 
when  the  OH  ion  is  introduced  into  solutions  containing  Cu  ions,  owing 
to  the  formation  of  complex  organic  ions  (see  also  page  117).  By  boil- 
ing with  the  simple  sugars,  the  cupric  portion  of  these  ions  is  reduced 
to  the  cuprous  condition  and  decomposition  into  the  simple  (Cu)2  ion 
takes  place.  In  the  presence  of  OH  ions,  (CuOH)2  forms  but  it  is  de- 
hydrated immediately  and  red  Cu20  is  precipitated.  This  reaction  is 

7 


82  QUALITATIVE   ANALYSIS. 

used  in  estimating  sugar  in  solution.     (Fehling's  Test.)      Cane  sugar 
must  be  inverted  before  it  will  cause  reduction. 

The  carbonate  ion,  C03,  from  alkaline  carbonate  solutions,  causes  a 
precipitation  of  a  mixture  of  hydroxid  and  carbonate  as  a  light  blue  pre- 
cipitate. Since  the  composition  of  this  substance  is  fairly  constant,  it  is 
probable  that  a  true  basic  salt  is  formed  of  the  nature  indicated  by  the 
formula : 

OH 

Cu 

C03 

/ 

Cu 

OH 
(See  page  21.) 

The  sulfid  ion,  S,  precipitates  black  CuS,  insoluble  in  dilute  acids 
and  alkalies.  Fairly  strong  HN03,  when  warm,  dissolves  it  owing  to 
oxidation  to  S  (or  H2SO4)  and  Cu(N03)2.  Copper  sulfid  is  somewhat 
soluble  in  yellow  ammonium  sulfid,  (NH4)2S.  xS,  less  in  yellow  sodium 
sulfid,  Na2S. xS,  and  not  soluble  in  colorless  Na2S  (difference  from  HgS 
and  Sb2S3  or  Sb2S5.  See  pages  108,  167  and  168.)  It  is  soluble  in 
alkaline  cyanids,  forming  soluble  complex  cyanids. 

The  sulfate  ion,  S04,  produces  no  precipitate,  for  the  sulfate  is  solu- 
ble in  about  4  parts  H20  at  ordinary  temperature's  and  less  than  2  parts 
at  100°.  From  solutions,  it  crystallizes  ordinarily  in  the  form  CuS04. 
5H20,  but  can  be  made  to  take  up  two  additional  molecules  and  conform 
to  the  type  of  the  "vitriols"  (see  page  188).  The  ordinary  salt  is 
known  commercially  as  * '  blue  vitriol. "  Its  blue  color  is  lost  when  it  is 
heated,  for  the  anhydrous  sulfate  is  dirty- white  in  color. 

The  chromate  ion,  Cr04,  precipitates  a  yellow  oxychromate,  having 
the  proportions  represented  by  the  formula,  3CuO.Cr03.2H20. 

The  chlorid  ion,  01,  causes  no  precipitation.  The  anhydrous  salt  is 
brown  but  it  separates  from  solution  in  the  form  of  green  crystals  hav- 
ing the  composition  CuCl2.2H20.  This  salt  is  deliquescent  and  dissolves 
in  less  than  its  own  weight  of  water. 

The  iodid  ion,  I,  probably  forms  cupric  iodid  at  first,  but  this  salt 
immediately  breaks  down  into  the  insoluble,  white  cuprous  iodid,  (Cul)2, 
with  the  separation  of  free  iodin,  which  colors  the  solution  brown. 

2CuI2-»(CuI)2  +  2I 


CUPKOUS    SALTS.  83 

/ 

The  cyanid  ion,  ON,  forms  the  very  unstable,  yellow  cupric  cyanid 
at  first,  but  this  decomposes  into  the  insoluble,  white  (CuCN)2  and  sets 
free  CN.  An  excess  of  the  cyanid  causes  the  precipitate  to  dissolve, 
owing  to  the  formation  of  complex  cyanids  (see  reaction  of  CN  ion  on 
(Cu)2  ion,  page  84). 

2Cu(CN)2-+  (CuCN)2  +  2CN 

The  ferrocyanid  ion,  (Fe  =  D6N6),  precipitates  gelatinous,  mahogany- 
colored  Cu2(Fe  =  C6N6). 

The  Cuprous  Ion,  (Cu)2. 

Cuprous  compounds  may  be  considered  as  containing  the  element  in 
the  univalent  condition.  All  of  the  compounds  seem  to  have  double 
molecules,  as  the  formulae  indicate.  (Note  similar  doubled  molecules 
of  mercurous  compounds,  page  110.) 

All  compounds  formed  by  the  cuprous  ion  are  so  insoluble  that  no 
solution  containing  an  appreciable  concentration  of  (Cu)2  ions  can  be 
made.  The  cuprous  halids  dissolve  readily  in  their  respective  acids, 
however,  forming  complex  cuprous-halogen  acids  of  the  type  H2CuCl3, 
H2CuBr3,  H2CuI3,  which  dissociate  into  the  complex  ions  CuCl3?  CuBr3 
and  CuI3.  If  such  a  solution  be  diluted,  the  complex  ions  decompose 
into  simple  ions  and  the  colorless  salts,  (CuCl)2,  (CuBr)2  and  (Cul)2, 
are  precipitated.  They  will  dissolve  in  NH4OH,  forming  complex  am- 
monia salts  which,  like  most  cuprous  compounds,  are  colorless  and  give 
colorless  solutions.  Oxygen  is  quickly  absorbed,  however,  and  the  deep 
blue  color  of  the  cupric-ammonia  compounds  appears  at  once. 

A  solution  of  a  cuprous  halid  containing  only  sufficient  acid  to  pre- 
vent precipitation  of  the  salt  will  serve  for  the  following  reactions. 

The  hydroxyl  ion,  OH,  forms  yellow  cuprous  hydroxid  at  first.  At 
ordinary  temperatures,  this  slowly  dehydrates  and  precipitates  red  Cu20. 
The  reaction  is  hastened  by  boiling. 

The  carbonate  ion,  C03,  produces  the  same  precipitate.  Since  no 
cuprous  salts  of  acids  containing  oxygen  are  known,  the  OH  ion,  pro- 
duced by  hydrolysis  of  the  carbonate,  is  the  active  agent  and  precipi- 
tates (CuOH)2. 

The  sulfid  ion,  S,  precipitates  black  Cu2S,  insoluble  in  alkalies  and  in 
dilute  acid  with  the  exception  of  warm  HN03,  which  oxidizes  it  to 
Cu(NO3)2  and  separates  free  sulfur. 

The  cyanid  ion,  CN,  precipitates  white  (CuCN)2,  soluble  in  excess 
of  alkaline  cyanids,  forming  complex  cyanids  of  the  proportions  indi- 


84  QUALITATIVE    ANALYSIS. 

cated  by  the  formulae  K3(CuC4N4)  and  K(CuC2N2).  In  fairly  concen- 
trated solutions,  the  ion  (CuC4N4)  is  very  stable,  decomposing  into  the 
simple  (Cu)2  ion  to  such  a  small  extent  that  the  S  ion  gives  no  precipi- 
tate. In  dilute  solutions,  this  ion  decomposes  almost  entirely  into 
(CuC2N2),  which  is  much  less  stable,  and  the  solution  is  precipitated  by 
the  S  ion. 

SILVER   (Argentum).     Ag.  107.9. 

Silver  is  a  brilliant  white  metal,  under  ordinary  conditions  not  at- 
tacked by  oxygen  or  other  constituents  of  the  atmosphere  at  any  tem- 
perature. It  unites  readily  with  the  halogens  and  with  sulfur,  with  the 
latter  element  forming  black  Ag2S.  Slight  traces  of  H2S  in  the  air 
account  for  the  tarnishing  of  silver  according  to  the  equation: 

4Ag  +  2H2S  +  02  ->  2Ag2S  +  2H20 

Metallic  silver  is  insoluble  in  HC1  but  readily  dissolves  in  HN03; 
also  in  hot  concentrated  H2SO4  with  the  evolution  of  S02.  In  the 
presence  of  oxygen,  silver  dissolves  in  solutions  of  the  alkaline  cyanids 
according  to  the  reaction: 

4Ag  +  8KCN  +  02  +  2H20  -»  4K(AgC2N2)  +  4KOH 

Silver,  like  copper,  shows  a  decided  disinclination  to  pass  into  the 
ionic  condition,  as  is  indicated  by  the  fact  that  considerable  energy  is 
absorbed  when  the  metal  becomes  the  ion.  This  partially  accounts  for 
the  fact  that  silver  does  not  dissolve  readily  in  acids  but  requires  also 
an  oxidizing  action  to  make  it  pass  into  the  ionic  form.  Silver  stands 
very  near  the  bottom  of  the  Potential  Series.  (See  page  327.)  If  metallic 
zinc,  iron,  lead,  tin,  mercury,  or  almost  any  other  metal  with  the 
exception  of  gold  and  platinum,  be  introduced  into  a  solution  contain- 
ing the  Ag  ion,  metallic  silver  is  deposited  while  the  other  metal  passes 
into  the  ionic  condition  and  goes  into  solution.  The  following  equation 
illustrates  the  reaction  and  shows  the  greater  tendency  of  these  metals 
to  pass  into  the  ionic  condition. 

Hg  +  2Ag  +  2N03  ->  Hg  +  2Ag  +  2NO3 

On  the  other  hand,  silver  passes  with  the  greatest  readiness  from  the 
ionic  into  the  metallic  condition  and  the  metal  may  be  obtained  from 
its  salts  without  difficulty  by  the  action  of  heat  and  reducing  agents. 

Heated  before  the  blowpipe  with  Na2C03  on  charcoal,  all  the  salts 
of  silver  are  decomposed.  A  brilliant  white  metallic  globule  is  formed, 


SILVER.  85 

without  coating  on  the  charcoal  (difference  from  Bi,  Pb  and  Sn),  which 
dissolves  readily  in  HN03  (difference  from  Sn)  and  gives  an  immediate 
precipitate  even  in  hot  dilute  solution  when  treated  with  HC1  (difference 
from  Pb  and  Bi). 

The  common  silver  ion  is  Ag.     In  the  presence  of  NH3,  complex 

+ 

silver-ammonia  ions  are  formed,  of  which  Ag(NH3)2  is  the  best  known. 
In  both  of  these  forms  silver  acts  as  kathion.  In  other  ions,  the  silver- 
cyanid,  AgC2N2,  and  a  very  complex  silver-thiosulfate  ion,  probably 
Ag2(S203)3,  the  metal  is  a  constituent  of  the  anion. 

Silver  compounds  are  usually  colorless  and  insoluble,  the  nitrate  being 
the  only  common  salt  which  is  readily  soluble.  It  dissolves  in  half  its 
weight  of  water  and  crystallizes  from  solution  without  water  of  crystal- 
lization, a  tendency  exhibited  by  all  silver  salts.  The  fact  that  the 
salts  of  this  metal  do  not  hydrolyze  is  an  indication  that  silver  is  a  base- 
forming  element  of  considerable  strength.  The  nitrate,  acetate  and  sul- 
fate  are  fairly  soluble,  but  may  be  precipitated  from  concentrated 
solutions.  Compounds  formed  by  the  complex  silver  ions  are  generally 

soluble. 

+ 

The  hydroxyl  ion,  OH,  added  to  solutions  containing  the  Ag  ion, 
causes  the  formation  of  AgOH,  which,  at  ordinary  temperatures,  im- 
mediately loses  water  and  forms  Ag20,  a  brown-black  precipitate,  soluble 
in  about  15,000  parts  of  water.  Its  solution  has  an  alkaline  reaction  and 
therefore  must  contain  OH  ions.  Hence  some  AgOH  must  exist  in  solu- 
tion. Measurements  of  its  electrical  conductivity  seem  to  indicate  that 
it  is  about  one-third  dissociated  at  this  dilution.  Consequently,  it  is  a 
very  much  weaker  base  than  the  fixed  alkalies  and  alkaline-earth  hy- 
droxids,  but  is  stronger  than  ammonium  hydroxid. 

Silver  oxid  is  entirely  decomposed  at  250°  into  oxygen  and  silver. 

When  precipitated  by  a  solution  of  |NaOH  or  KOH,  silver  oxid  is 
not  affected  by  an  excess  of  the  reagent.  If  NH4OH  is  sparingly  added 
to  a  solution  containing  Ag  ions,  Ag2O  is  precipitated,  but  dissolves  in 
the  slightest  excess  of  the  precipitant.  Owing  to  the  decomposition  of 
the  undissociated  NH4OH  into  H2O  and  NH3,  a  complex  ion,  Ag(NH3)2, 
is  formed,  the  salts  of  which  are  soluble.  Since  this  complex  ion  forms 
with  great  readiness  and  is  decomposed  but  little  into  NH3  and  the 
simple  silver  ion,  the  presence  of  NH3  decreases  the  concentration  of 
the  Ag  ion  very  decidedly  (see  page  36),  hence  the  precipitated  Ag2O, 
and,  similarly,  most  other  salts  of  silver,  dissolve  readily  in  NH4OH. 

The  silver-ammonia  ion  and,  less  readily,  the  simple  Ag  ion,  may  be 
reduced  with  the  greatest  ease  and  metallic  silver  may  be  deposited  by  the 


86  QUALITATIVE    ANALYSIS. 

action  of  carbonaceous  matter,  in  general,  and,  especially,  by  tartaric 
and  citric  acids,  aldehyde  and  the  sugars.  Inasmuch  as  part  of  the 
silver  adheres  closely  to  the  glass  vessel  in  which  the  reduction  takes 
place  and  forms  a  brilliant  mirror,  this  reaction  is  used  as  a  test  for 
certain  of  these  substances.  (See  page  285.) 

The  stains  which  are  caused  by  silver  solutions  coming  in  contact 
with  organic  matter  are  due  to  this  same  reducing  action,  the  finely  di- 
vided metallic  silver  being  black  in  color.  A  solution  of  potassium  cyanid, 
by  dissolving  the  metal,  will  remove  the  stain. 

The  carbonate  ion,  C03,  causes  the  precipitation  of  white  Ag2C03, 
which  soon  undergoes  partial  decomposition  into  C02  and  Ag20,  the 
latter  giving  a  yellow  tinge  to  the  precipitate.  If  (NH4)2C03  is  used  as 
the  precipitant,  Ag2C03  dissolves  somewhat  in  excess,  since  (NH4)2COS 
hydrolyzes  considerably.  Free  NH3  is  therefore  present  in  the  solution 
and  causes  the  formation  of  the  soluble  silver-ammonia  ion. 

The  sulfid  ion,  S,  precipitates  black  Ag2S  from  both  acid  and  alka- 
line solutions.  This  compound  dissolves  when  treated  with  HN03  since 
it  is  oxidized  to  free  sulfur  (or  H2S04)  and  AgNO3.  Silver  sulfid  is  an 
extremely  insoluble  substance,  hence  it  does  not  dissolve  appreciably  even 
in  the  strong  acids,  nor  in  NH4OH  or  dilute  KCN. 

The  sulfate  ion,  S04,  precipitates  white  Ag2S04  from  neutral  solu- 
tions only,  since  it  is  exceedingly  soluble  in  traces  of  acid  and  of 
ammonia. 

The  chromate  ion,  Cr04,  precipitates  red  Ag2Cr04,  which,  being  the 
salt  of  an  acid  only  moderately  strong,  is  soluble  in  the  strong  acids. 
Like  other  silver  salts  which  are  appreciably  soluble,  it  dissolves  in 
NH4OH  with  the  formation  of  the  silver-ammonia  ion. 

The  chlorid  ion,  01,  precipitates  curdy  white  AgCl,  soluble  in  about 
700,000  parts  of  water,  insoluble  in  all  dilute  acids.  It  dissolves  with 
extreme  readiness  in  NH4OH,'  also  in  KCN  and  Na2S203.  In  the  pres- 
ence of  the  cyanid  ion,  CN,  the  silver  ion  forms  the  complex  silver-cyanid 
ion,  Ag(CN)2. 

Ag  +  2CN-»AgC2N2 

Since  AgC2N2  is  very  stable  and  decomposes  but  little  into  the  simple  ions, 
the  presence  of  the  CN  ion  reduces  the  concentration  of  Ag  ions  de- 
cidedly and,  therefore,  the  AgCl  dissolves  readily.  (See  page  36.) 

In  sodium  thiosulfate,  AgCl  dissolves  forming  a  solution  from  which 
crystallizes  a  salt  having  the  proportions  represented  by  the  formula 
(Na2S2O3)2.Ag2S2O3.  The  complex  ion  formed  by  this  salt  is  very  stable, 


SILVER    HALIDS.  87 

but,  when  boiled  with  acids,  decomposes  with  the  formation  of  black  Ag2S. 

In  concentrated  solutions  of  HC1  and  other  chlorids,  AgCl  dissolves 
readily,  probably  forming  complex  compounds  of  the  type  HAgCl2, 
which  separate  Ag  ions  only  on  diluting,  in  which  case  AgCl  is  again 
precipitated. 

The  bromid  ion,  Br,  precipitates  white  AgBr,  soluble  in  about 
1,000,000  parts  of  water,  and  very  similar  to  AgCl.  It  dissolves  in  KCN 
and  Na2S203  as  does  AgCl  and  for  similar  reasons.  It  is  less  readily 
soluble  in  NH4OH,  however. 

The  iodid  ion,  I,  precipitates  pale  yellow  Agl  which  resembles  the 
'other  halids  of  silver  and  is  soluble  in  about  1,500,000  parts  of  water. 
Like  them  it  dissolves  readily  in  KCN  and  Na2S2O3,  but  is  not  ap- 
preciably soluble  in  NH4OH.  The  cause  ior  this  varying  solubility  of 
the  silver  halids  in  NH4OH  will  be  apparent  from  the  following 

reasoning. 

+ 
In  a  saturated  solution  of  AgCl  there  will  be  concentration  of  Ag 

ions  which  may  be  represented,  let  us  say,  by  100.  Since  AgBr  is  less 
soluble  than  AgCl,  the  concentration  of  Ag  ions  in  a  saturated  solution 
of  the  latter  salt  will  be  less,  and  may  be  represented  by  75.  Being  less 

soluble  than  either,  Agl  will  give  a  concentration  which,  on  the  same 

+ 
scale,  may  be  taken  as  50.     The  concentration  of  Ag  ions  due  to  the 

decomposition  of  the  very  soluble-but-stable,  complex  silver-ammonia  ion 

+ 

may  be  represented  by  35.  The  same  ion,  Ag(NH3)2,  is  formed  by  dis- 
solving AgCl,  AgBr  and  Agl  in  NH4OH.  Therefore  the  final  concen- 

+ 

tration  of  Ag  ions  will  be  the  same  in  each  case.  But  the  relative  reduc- 
tion in  the  case  of  the  chlorid  is  much  greater  than  with  the  bromid, 
therefore,  AgCl  is  more  soluble  in  NH4OH  than  AgBr  is.  The  reduction 
in  the  case  of  the  iodid  is  relatively  too  small  to  cause  appreciable  solu- 
tion. The  fact  that  all  three  silver  halids  dissolve  in  KCN  and  Na2S2O3 
is  due  to  a  stability  on  the  part  of  the  silver-cyanid  and  silver-thiosulfate 
ions  greater  than  that  possessed  by  the  silver-ammonia  ion.  Therefore 
the  final  concentration  of  Ag  ions  is  less  in  every  case  and  the  relative 
reduction  is  sufficiently  great  to  cause  solution. 

On  the  addition  of  acids  to  a  solution  containing  the  silver-ammonia 
ion,  a  process,  the  reverse  of  that  by  which  it  is  formed,  takes  place. 

4-  4- 

The  H  ions  combine  with  NH3  to  form  NH4  ions,  thus  causing  the 
complex  ion  to  decompose  and  give  rise  to  the  simple  Ag  ion.  In  this 

4- 

manner,  the  concentration  of  the  Ag  ion  is  increased  and  AgCl,  AgBr, 
or  Agl  is  precipitated  once  more. 

By  exposure  to  light  all  of  the  silver  halids  turn  dark  and  liberate 


88 


QUALITATIVE   ANALYSIS. 


a  part  of  the  halogen,  forming  probably  a  compound  of  a  lower  type 
such  as  Ag2CL  The  original  color  of  a  darkened  silver  halid  is  restored 
by  chlorin  or  bromin  water. 

All  the  silver  halids  may  be  decomposed  by  bringing  them  in  con- 
tact with  metallic  zinc  and  covering  them  with  H2S04.  By  the  action  of 
the  nascent  hydrogen  formed  by  the  interaction  of  the  acid  and  the 
metal,  metallic  silver  is  thrown  out  and  hydrogen-halogen  acids  formed. 

The  cyanid  ion,  CN,  precipitates  white  AgCN,  insoluble  in  HN03, 
but  readily  soluble  in  NH4OH,  Na2S203,  and  an  excess  of  alkaline  cyanid. 

The  phosphate  ion,  ^P04,  precipitates  yellow  Ag3P04. 

The  arsenate  ion,  As04,  separates  chocolate-colored  Ag3As04. 

The  arsenite  ion,  As03,  causes  yellow  Ag3As03  to  fall. 

As  would  be  expected,  all  of  these  salts  are  soluble  in  HN03  and 
NH4OH  as  well,  and  can  be  precipitated  only  from  neutral  solutions. 

+ 

Inasmuch  as  the  Ag  ion  is  an  important  reagent  in  detecting  and 

differentiating  between  acid  radicles  or  anions,  the  following  is  a  list  of 

+ 

the  salts  formed  by  the  Ag  ion  with  those  acids  which  will  be  considered 
in  the  present  scheme  of  analysis : 


Acid. 

Salt. 

Color. 

H20. 

HNOS. 

NH.OH. 

Acetic 

Ag(C2H302) 

White 

Sol.* 

Sol. 

Sol. 

Arsenic 

Ag3AsO4  ' 

Chocolate 

Insol. 

Sol. 

Sol. 

Arsenious 

Ag3AsO3 

Yellow 

Insol. 

Sol. 

Sol. 

Boric 

AgBO2 

White 

Insol. 

Sol. 

Sol. 

Carbonic 

Ag2C03 

White 

Insol. 

Sol. 

Sol. 

Chloric 

AgC103 

White 

Sol. 

Sol. 

Sol. 

Chromic 

Ag2Cr04 

Bed 

Insol. 

Sol. 

Sol. 

Cyanic 

AgOCN 

White 

Insol. 

Sol. 

Sol. 

Ferricyanic 

Ag3(FeC6N6) 

Orange 

Insol. 

Sol. 

Sol. 

Ferrocyanic 

Ag4(FeCGN.) 

White 

Insol. 

Sol. 

Insol. 

Hydrobromic 

AgBr 

White 

Insol. 

Insol. 

Sol. 

Hydrochloric 

AgCl 

White 

Insol. 

Insol. 

Sol. 

Hydrocyanic 

AgCN 

White 

Insol. 

Insol. 

Sol. 

Hydrofluoric 

AgF 

White 

Sol. 

Sol. 

Sol. 

Hydriodic 

Agl 

Yellow 

Insol. 

Insol. 

Insol. 

Hydrosilicofluoric 

Ag2SiF6 

White 

Sol. 

Sol. 

Sol. 

Hydrosulfuric 

Ag2S 

Black 

Insol. 

Sol. 

Insol. 

Hypochlorous 

AgOCl  (unstable) 







Nitric 

AgNO3 

White 

Sol. 

Sol. 

Sol. 

Nitrous 

AgNO2 

White 

Sol.* 

Sol. 

Sol. 

Oxalic 

Ag2C204 

White 

Insol. 

Sol. 

Sol. 

Perchloric 

AgClO4 

White 

Sol. 

Sol. 

Sol. 

Permanganic 
Phosphoric 

AgMnO4 
Ag3P04 

Purple 
Yellow 

Sol. 
Insol. 

Sol. 
Sol. 

Sol. 
Sol. 

Silicic 

(unknown) 

.  







Sulfuric 

Ag2S04 

White 

Sol.* 

Sol. 

Sol. 

Sulfurous 

Ag2S03 

White 

Insol. 

Sol. 

Sol. 

Thiocyanic 

AgSCN 

White 

Insol. 

Insol. 

Sol. 

Thiosulfuric 

Ag2S203 

White 

Insol. 

Sol. 

Sol. 

Tartaric 

Ag2(C4H404) 

White 

Insol. 

Sol. 

Sol. 

May  be  precipitated  from  fairly  concentrated  solutions. 


GOLD.  89 

GOLD   (Aurum).    Au  — 197.2. 

Gold  is  a  soft,  malleable,  heavy,  yellow  metal.  It  is  not  acted  upon 
by  the  constituents  of  the  atmosphere  at  any  temperature  and,  with  sil- 
ver, platinum  and  a  few  other  elements,  is  known  as  a  ' '  noble  metal. ' ' 

Gold  does  not  dissolve  in  any  acid  but,  since  it  is  attacked  by  the 
halogens  at  ordinary  temperatures,  it  dissolves  readily  in  aqua  regia 
because  of  the  free  chlorin  which  this  reagent  contains.  Like  the  other 
'  *  noble  metals, ' '  gold  shows  a  very  strong  disinclination  to  pass  into  the 
ionic  condition  and  energy  is  absorbed  when  the  metal  takes  on  the  elec- 
tric charge.  Gold  stands  at  the  very  bottom  of  the  Potential  Series, 
hence  if  any  other  metal  is  introduced  into  a  solution  containing  salts 
of  gold,  the  latter  metal  is  deposited  and  the  former  goes  into  solution. 

All  gold  salts  are  unstable  and  are  readily  reduced  to  the  metallic 
condition  by  heat  and  by  reducing  agents  of  all  kinds.  Ferrous  and 
stannous  salts,  oxalic  and  sulfurous  acids  all  effect  the  reduction  to  the 
metal  in  acid  solutions  at  ordinary  temperatures.  When  stannous 
chlorid  is  used  as  a  reducing  agent,  the  finely  divided  gold  is  of  a  red 
or  purple  color  and  under  the  name  of  "purple  of  Cassius"  affords  a 
well-known  test  for  this  metal. 

Heated  before  the  blowpipe  on  charcoal  with  Na2C03,  all  gold  salts 
leave  a  yellow  globule  of  the  metal,  insoluble  in  nitric  acid. 

With  oxygen,  metallic  gold  does  not  unite  directly,  nevertheless  by 
indirect  means  two  oxids  may  be  obtained,  a  dark  violet  monoxid,  Au20, 
and  a  brown  trioxid,  Au203.  Each  of  these  substances  may  be  consid- 
ered as  the  basis  from  which  a  series  of  gold  salts  is  derived. 

The  univalent  aurous  salts  are  nearly  colorless,  generally  insoluble 
and  are  decomposed  slowly  by  water  according  to  the  equation: 

3AuCl  ->  AuCl3  +  2Au 

For  these  reasons  the  aurous  ion,  Au,  is  but  little  known. 

Aurous  cyanid,  AuCN,  unites  with  potassium  cyanid  to  form  a  com- 
pound, KAuC2N2,  in  which  the  gold  is  a  constituent  of  the  anion, 
AuC2N2.  This  compound  is  fairly  stable  and  is  of  importance  in  the 
metallurgy  of  gold,  for,  in  the  presence  of  air,  metallic  gold  dissolves  in 
potassium  cyanid  solutions  according  to  the  equation: 

4Au  +  8KCN  +  02  +  H20  ->  4KAuC2N2  +  4KOH 

Auric  salts  are  trivalent  and  somewhat  more  stable  than  aurous  salts. 
They  are  generally  light  yellow  in  color,  usually  soluble  and  in  solution 
give  rise  to  the  auric  ion,  Au. 


90  QUALITATIVE    ANALYSIS. 

Gold  salts  of  acids  containing  oxygen  are  but  little  known  since  they 
are  very  unstable  and  decompose  readily  even  in  solution. 

Auric  chlorid,  AuCl3,  is  the  most  important  gold  salt.  It  is  a  light- 
yellow,  deliquescent  substance  which  at  180°  breaks  down  into  aurous 
chlorid  and  free  halogen  and,  at  a  somewhat  higher  temperature,  is 
entirely  decomposed  into  its  elements.  In  solution,  it  gives  rise  to  the 

+-H- 

auric  ion,  Au.  Auric  chlorid  readily  unites  with  hydrochloric  acid  to 
form  a  crystalline  compound,  hydrochlorauric_  acid,  HAuCl4,  in  which 
the  gold  acts  as  a  constituent  of  the  anion,  AuCl4.  Several  salts  of  this 
acid  are  known.  Of  these  the  sodium  compound,  NaAuCl4,  finds  use  in 
photography  as  a  "toning  solution." 

The  hydroxyl  ion,  OH,  added  to  a  solution  containing  the  auric  ion, 

-H-f- 

Au,  causes  a  red,  flocculent  precipitate  which  seems  almost  identical  in 
appearance  with  ferric  hydroxid.  Like  aluminum,  zinc,  lead  and  tin 
and  certain  other  hydroxids,  auric  hydroxid  acts  either  as  an  acid  or  as 
a  base.  It  is  readily  soluble  in  acids,  forming  salts  in  which  gold  plays 
the  part  of  a  metal.  It  dissolves  in  alkalies  as  well,  giving  rise  to  com- 
pounds of  the  composition  KAu02.  These  salts  are  known  as  aurates, 
for  in  them  gold  acts  as  an  acid-forming  element.  So  readily  does  auric 
hydroxid  dissolve  in  an  excess  of  sodium  or  potassium  hydroxid  that,  in 
dilute  gold  solutions,  no  precipitate  is  ordinarily  obtained. 

Ammonium  hydroxid,  in  distinction  from  sodium  and  potassium 
hydroxids,  precipitates  from  solutions  containing  the  Au  ion,  a  floccu- 
lent, brown  powder  which,  when  dry,  is  explosive.  It  has  the  composi- 
tion, AuN2H3.3H20,  and  is  known  as  "fulminating  gold." 

The  carbonate  ion,  C03,  from  solutions  of  alkaline  carbonates,  be- 
cause of  hydrolysis  acts  as  does  the  hydroxid  ion. 

The  sulfid  ion,  S,  precipitates  a  black  substance,  the  composition  of 
which  varies  from  Au2S3  to  Au2S  or,  apparently,  it  may  contain  metallic 
gold.  The  substance  is  insoluble  in  acids,  but  dissolves  in  aqua  regia 
and  in  solutions  of  alkaline  sulfids,  forming  in  the  latter'  case  a  thioaurate 
analogous  to  thioantimonates  and  similar  compounds.  This  fact  causes 
gold  to  be  separated  in  qualitative  analysis  along  with  antimony  and 
arsenic  in  the  tin  group. 

Group  2,  First  Type. 

Beryllium  9.1 ;       Magnesium  24.4 ;       Calcium  40.1 ;       Strontium  87.6 ; 
Barium  137.4;       Radium  225. 

The  alkaline  earths,  as  this  group  is  called,  are  soft,  light,  lustrous 
metals.  Although  they  are  unstable  in  the  air  and  decompose  water,  the 


UN1VE        ^THE*  ALKALINE    EAETH    GROUP. 


91 

action  i?fiS^Wr^i§orous  as  with  the  alkali  metals.  The  chemical  activity 
of  the  elements  increases  as  the  atomic  weight  rises,  but  all  these  sub- 
stances are  reactive,  forming  stable  compounds  from  which  the  elements 
may  be  obtained  only  with  difficulty.  The  members  of  this  group  always 
show  a  valence  of  two.  The  oxids  unite  readily  with  water  to  form  hy- 
droxids  which  act  as  moderately  strong  bases,  barium  hydroxid  disso- 
ciating somewhat  more  than  magnesium  hydroxid.  Salts  of  this  group 
are  generally  colorless.  The  solubility  of  the  hydroxids  increases,  and 
of  the  sulfates  decreases,  as  the  atomic  weight  rises.  With  the  excep- 
tion of  magnesium,  salts  of  the  common  metals  give  color  to  a  Bunsen 
flame. 

Beryllium,  and  radium  especially,  are  very  rare  elements,  and  they 
will  not  be  considered  in  detail. 

MAGNESIUM.    Mg  — 24.4. 

Magnesium  is  a  light,  silver- white,  crystalline  metal.  At  ordinary 
temperatures  it  tarnishes  somewhat  in  the  atmosphere,  forming  a  com- 
pact film  of  the  oxid  which  acts  as  a  protective  coating.  Heated  in  the 
air,  it  burns  with  an  intense  white  light  to  the  oxid.  Because  of  the 
great  actinic  power  of  this  light,  it  is  much  used  in  photography  for 
"flash  lights."  The  metal  is  capable  of  taking  nitrogen  as  well  as 
oxygen  from  the  air,  forming  the  nitrid,  Mg3N2. 

The  metal  passes  readily  into  the  ionic  condition  and,  therefore, 
dissolves  in  all  acids.  Even  the  concentration  of  H  ions  due  to  the  dis- 
sociation of  water  is  sufficient  at  100°  to  liberate  hydrogen  and  form 
Mg(OH)2.  This  substance,  being  insoluble,  soon  coats  over  the  metal 
and  prevents  further  action.  If  ammonium  salts  be  present  in  solu- 
tion to  dissolve  the  hydroxid  (see  later),  the  action  proceeds  freely. 
Magnesium  is  not  precipitated  from  solutions  of  its  salts  by  other  metals, 
but,  on  the  contrary,  with  the  exception  of  aluminum  and  chromium, 
under  proper  conditions  magnesium  will  precipitate  all  the  common 
metals  from  solutions  of  their  salts,  provided  the  metals  do  not  react  with 
water.  When  magnesium  passes  into  the  ionic  condition,  more  energy  is 
liberated  than  in  the  case  of  any  other  metal  with  the  exception  of  the 
alkali  and  alkaline-earth  metals.  For  this  reason  its  salts  are  reduced  in 
the  "dry  way"  only  and  with  considerable  difficulty. 

Magnesium  oxid,  MgO,  is  called  ' '  magnesia, ' '  and  like  oxids  generally 
it  may  be  formed  by  burning  the  metal  or  by  heating  the  hydroxid,  car- 
bonate, nitrate  or  salts  of  organic  acids  (which  break  down  first  into 
carbonates).  It  fuses  only  at  the  highest  obtainable  temperatures.  It 


92  QUALITATIVE    ANALYSIS. 

is  a  light  white  substance,  but  slightly  soluble  in  water,  reacting  with  it 
to  form  the  hydroxid. 

Magnesium  salts  are  generally  colorless  and  usually  soluble.  Of 
the  common  salts,  only  the  hydroxid,  carbonate  and  phosphate  are  in- 
soluble. Magnesium  salts  contain  the  metal  in  the  bivalent  condition, 
acting  as  a  base-forming  element  distinctly  stronger  than  the  ordinary 
metals,  yet  not  sufficiently  strong  to  protect  its  compounds  from  hydrol- 
ysis. They  give  no  color  to  a  Bunsen  flame. 

The  simple  magnesium  ion  is  Mg.    Some  complex  ions  are  known. 

For  the  reactions  of  the  Mg  ion,  any  soluble  simple  salt  will  serve. 
The  nitrate,  Mg(N03)2.6H2O,  is  deliquescent  and  soluble  in  half  its 
weight  of  water.  The  acetate,  Mg(C2H302)2.4H2O,  is  also  deliquescent 
and  extremely  soluble. 

The  hydroxyl  ion,  OH,  precipitates  white,  gelatinous  Mg(OH)2,  sol- 
uble in  about  100,000  parts  of  pure  water,  but  the  presence  of  small 
amounts  of  impurities  generally  increases  its  solubility  greatly.  It  dis- 
sociates to  nearly  the  same  extent  as  the  other  hydroxids  of  this  group 
at  the  same  dilution,  but,  owing  to  its  insolubility,  affects  litmus  slowly. 

If  NH4OH  be  used  as  the  reagent  for  introducing  the  OH  ion,  only  a 
partial  precipitation  of  the  magnesium  as  hydroxid  is  effected.  This 
phenomenon  is  due  partly  to  the  reduced  concentration  of  OH  ions 
caused  by  the  repression  of  the  ionization  of  the  weak  base,  NH4OH,  by 
its  more  readily-dissociating  salts  (see  page  31),  and  partly  owing  to  the 
formation  of  soluble,  complex  magnesium-ammonium  salts  of  the  type 

-H- 

(NH4)2MgCl4.  These  compounds  do  not  separate  the  simple  Mg  ion  at 
ordinary  dilutions  to  any  great  extent,  but  they  dissociate  into  (NH4) 
and  MgCl4  ions.  In  the  presence  of  sufficient  NH4C1  (or  other  am- 
monium salt),  no  precipitation  of  Mg(OH)2  will  take  place.  A  solution 
of  this  character  (usually  1  gram-molecule  of  MgCl2  plus  2  of  NH4C1 
rendered  strongly  alkaline  with  NH4OH)  is  known  as  "magnesia  mix- 
ture "  and  is  used  as  a  precipitant  for  phosphoric  and  arsenic  acids. 
This  same  property  of  forming  double  salts  with  ammonium  compounds 
is  possessed  by  cobalt,  nickel,  manganese  zinc  and  ferrous  iron  as  well. 

The  carbonate  ion,  C03,  causes  a  white  precipitate,  consisting  of  a 
mixture  of  MgCO3  and  Mg(OH)2,  owing  to  hydrolysis.  If  (NH4)2CO3 
be  used  as  the  precipitant,  only  a  partial  precipitation  ensues,  owing  to 
the  formation  of  the  complex  salt  already  mentioned.  In  the  presence 
of  ammonium  salts,  no  precipitate  will  fall. 

The  precipitate  obtained  with  alkaline  carbonates  dissolves  in  water 
containing  carbon  dioxid  (i.  e.,  carbonic  acid),  forming  a  soluble  acid 


MAGNESIUM.  93 

carbonate  similar  in  all  respects  to  the  corresponding  salt  of  calcium. 
(See  page  95.)  By  slow  evaporation  of  such  a  solution,  the  normal  car- 
bonate, MgC03,  which  is  the  mineral  magnesite,  crystallizes.  Like  cal- 
cium salts,  magnesium  salts  in  solution  make  water  "hard." 

The  sulfid  ion,  S,  causes  no  precipitate,  for  MgS  is  formed  only  in 
the  "dry  way"  and  in  aqueous  solution  is  hydrolyzed  into  Mg(OH)2  and 
H2S.  Since  Mg(OH)2  is  insoluble  and  is  thus  removed  from  the  sphere 
of  action,  complete  hydrolysis  takes  place,  the  reaction  proceeding  prac- 
tically only  in  the  direction  indicated  by  the  equation : 

MgS  +  2HOH  -»  Mg(OH)2  +  H2S 

If  (NH4)2S  is  used  to  introduce  the  S  ion,  a  partial  precipitation  of 
Mg(OH)2  may  be  obtained  if  no  other  ammonium  salts  are  present. 
With  H2S,  however,  no  Mg(OH)2  is  precipitated  because  of  the  acid 
formed  during  the  metathesis  indicated  by  the  equation: 

MgCl2  +  H2S  ->  MgS  +  2HC1 

The  sulfate  ion,  S04,  causes  no  precipitation  since  MgS04.7H20  is 
soluble  in  less  than  three  parts  of  H2O,  crystallizing  from  solutions  in 
the  form  characteristic  of  the  "vitriols."  (See  page  188.)  It  is  the  only 
common  salt  of  magnesium  that  does  not  deliquesce.  Under  the  name 
of  "Epsom  salt,"  it  finds  medicinal  use. 

The  chromate  ion,  Cr04,  causes  no  precipitate,  as  MgCr04  is  very 
soluble. 

The  chlorid  ion,  01,  causes  no  precipitate,  for  MgCl2.6H20,  like  the 
bromid,  iodid  and  cyanid,  is  extremely  soluble,  dissolving  in  half  its 
weight  of  water.  In  solution,  it  hydrolyzes  considerably  and,  on  evapo- 
ration, the  volatile  HC1  escapes,  leaving  behind  an  oxychlorid,  Mg2OCl2. 
For  this  reason  chiefly,  sea-water  can  not  be  used  in  boilers  as  the  acid 
corrodes  the  metal. 

The  fluorid  ion,  F,  precipitates  white  MgF2  which  is  only  slowly  soluble 
in  the  stronger  acids. 

The  phosphate  ion,  P04,  precipitates  white  Mg3(P04)2  which  is 
soluble  in  about  5000  parts  of  H2O.  If  HNa2PO4  is  used  as  the  precipi- 
tating reagent,  tertiary  magnesium  phosphate  does  not  fall,  but  rather  the 
secondary  or  acid  magnesium  phosphate,  HMgPO4,  which  is  soluble  in 
less  than  500  parts  of  H2O.  In  the  presence  of  NH4OH,  however, 
(NH4)MgP04.6H20  is  precipitated  which,  on  boiling,  assumes  a  crystal- 
line form.  This  salt  is  soluble  in  about  10,000  parts  of  H2O,  but,  in  the 


94  QUALITATIVE    ANALYSIS. 

presence  of  considerable  NH4OH,  requires  more  than  50,000  parts  for 
solution. 

The  oxalate  ion,  Cj04,  may  precipitate  from  solutions  containing  the 
Mg  ion  white  magnesium  oxalate,  MgC204.2H2O,  which  requires  1500 
parts  of  H2O  for  solution.  An  excess  of  any  alkaline  oxalate,  however, 
forms  double  oxalate  salts  of  several  different  proportions,  all  of  which 

are  soluble. 

CALCIUM.     Ca  —  40.1. 

Calcium  is  a  silvery-white,  crystalline,  malleable  metal.  In  dry  air 
it  is  stable  but,  under  ordinary  conditions,  it  is  quickly  attacked,  forming 
CaC03  and  liberating  hydrogen.  It  has  no  industrial  uses  and  is  merely 
of  scientific  interest. 

The  metal  dissolves  very  readily,  even  in  dilute  solutions  of  acids  as 
weak  as  H2C03,  forming  salts  in  which  the  element  is  always  divalent. 
The  concentration  of  H  ions,  given  by  pure  H20,  is  sufficient  to  cause 
the  metal  to  dissolve  slowly  at  ordinary  temperatures,  liberating  hydro- 
gen and  forming  the  hydroxid  according  to  the  reaction: 


When  the  element  passes  into  the  ionic  condition  a  large  amount  of 
energy  is  liberated,  hence  its  salts  are  extremely  stable.  The  metal  acts 
as  a  base-forming  element  of  considerable  strength,  consequently  salts  of 
the  strong  acids  show  no  hydrolysis.  With  few  exceptions  calcium  salts 
are  colorless. 

Calcium  salts  color  a  Bunsen  flame  orange.  This  indication  is  some- 
what masked  by  the  presence  of  other  members  of  this  group.  If  the 
salt  which  is  heated  be  moistened  with  HC1,  the  calcium  flame  will  flash 
out  for  a  moment  at  the  start,  since  CaCl2  requires  less  heat  to  volatilize 
and  decompose  it  than  either  BaCl2  or  SrCl2.  Barium  and  strontium 
will  appear  after  the  calcium  is  gone  and  the  indication  of  these  elements 
will  be  more  lasting. 

With  the  spectroscope,  calcium  is  detected  chiefly  by  a  green  and  an 
orange  line  about  equally  distant  from  the  ever-present  sodium  line. 
(See  frontispiece.) 

The  only  calcium  ion  definitely  known  is  Ca.  For  the  reactions  of  this 
ion  any  soluble  calcium  salt  will  serve.  The  nitrate,  Ca(N03)2.4H20  is 
deliquescent  and  soluble  in  about  one-quarter  of  its  weight  of  water. 
The  anhydrous  salt  is  readily  soluble  in  amyl  alcohol  and  by  this  means 
may  be  separated  from  barium  and  strontium,  which  are  insoluble  in  the 
same  solvent.  (See  page  270.) 


CALCIUM    SALTS.  95 

The  acetate,  Ca(C2H302)2.5H20,  is  efflorescent. 

The  hydroxyl  ion,  OH,  produces  no  precipitate  under  ordinary  con- 
ditions, for  Ca(OH)2  is  soluble  in  about  800  parts  of  cold  water.  It 
requires  more  than  1600  parts  at  100°,  for  its  solubility  decreases  as  the 
temperature  rises.  Concentrated  solutions  of  NaOH  or  KOH  may  pro- 
duce a  precipitate  but  the  OH  ion  is  never  present  in  NH4OH  solutions 
in  concentration  sufficient  to  do  so. 

Calcium  hydroxid  is  a  strong  base,  its  dissociation  being  surpassed 
in  a  marked  degree  only  by  the  hydroxids  of  the  alkali  metals.  As  it  is 
by  far  the  cheapest  reagent  of  this  kind,  it  is  used  for  all  industrial  pur- 
poses where  a  strong  base  is  required.  It  is  obtained  by  "burning  lime- 
stone," i.  e.,  heating  the  carbonate,  CaC03,  which  loses  C02  and  forms 
the  oxid,  CaO,  which  is  lime.  Treated  with  water,  lime  "slacks"  and 
forms  the  hydroxid  or  slacked  lime.  Like  all  strong  bases,  calcium 
hydroxid  absorbs  C02  from  the  atmosphere,  thus  reforming  the  carbon- 
ate. A  solution  of  Ca  (OH)  2  is  known  as  '  '  lime-water.  '  ' 

The  carbonate  ion,  C03,  precipitates  white,  amorphous  CaC03  which, 
on  warming,  readily  passes  into  the  crystalline  condition,  which  is  the 
mineral  calcite.  During  this  change,  liquids  containing  CaC03  in  sus- 
pension become  much  less  opaque  since  the  crystalline  form  is  much 
more  dense  and  compact.  On  standing,  it  settles  rapidly.  At  a  bright 
red  heat,  i.  e.,  about  800°,  CaC03  breaks  down  readily  into  C02  and  CaO. 

Calcium  carbonate  is  soluble  in  about  60,000  parts  of  H20.  It  is 
much  more  readily  soluble  in  water  containing  C02,  forming  acid  cal- 
cium carbonate,  which  dissolves  in  about  1000  parts  of  H20. 

H—  O 

C  =  0 
0  O 

/  \  / 

Ca          C  =  0  +  C02  +  H20-*Ca 

\  /  \ 

0  O 


H—  O 

On  heating,  all  acid  carbonates  are  decomposed  into  normal  carbon- 
ates, water  and  CO,,  thus  reversing  the  reaction.  Lime  salts,  together 
with  magnesium  salts,  present  in  solution  make  water  '  *  hard,  '  '  i.  e.,  form 


96  QUALITATIVE    ANALYSIS. 

insoluble  compounds  with  soap,  thus  preventing  its  cleansing  action. 
When  these  elements  are  present  as  acid  carbonates,  water  is  said  to  be 
"temporarily  hard,"  for,  on  boiling,  normal  carbonates  are  precipitated 
and  the  "hardness"  removed.  On  exposing  a  solution  of  acid  calcium 
carbonate  to  the  air,  the  same  reaction  takes  place  slowly  as  the  excess 
of  CO2  escapes.  In  this  way  CaC03  is  deposited  as  stalagmites  and 
stalactites  in  caves. 

In  the  presence  of  NH4C1,  CaC03  dissolves,  on  boiling,  owing  to  the 
metathesis — 

CaC03  +  2NH4C1-^  CaCl2  +  (NH4)2C03 

When  heated,  (NH4)2C03  decomposes  into  NH3,  C02  and  H20,  which 
are  removed  from  solution  by  boiling,  thus  causing  the  reaction  to  pro- 
ceed in  the  direction  indicated.  For  this  reason,  in  the  presence  of  con- 
siderable NH4C1,  a  small  amount  of  (NH4)2CO3  may  not  precipitate  cal- 
cium and  the  other  members  of  this  group  as  carbonates  completely, 
especially  if  the  solution  is  hot,  as  is  usually  necessary. 

The  sulfid  ion,  S,  under  ordinary  conditions  produces  no  precipitate 
for  CaS,  though  somewhat  insoluble,  is  hydrolyzed,  forming  soluble  prod- 
ucts, first  CaSH  and  finally  Ca(OH)2  and  H2S. 

The  sulfate  ion,  S04,  causes  a  white  precipitate  of  CaS04.2H20  in 
solutions  of  fair  concentration  only,  since  the  precipitate  is  soluble  in 
about  500  parts  of  water.  If  CaS04  in  solution  be  used  as  the  precipi- 
tating reagent,  obviously  no  precipitate  will  form  (distinction  from  Ba 
andSr). 

CaCl2  +  CaS04  -»  CaS04  +  CaCl2 

The  naturally-occurring  sulfate,  gypsum,  is  heated  until  it  has  lost 
most  of  its  water  of  crystallization.  The  dehydrated  product  is  known 
as  "plaster  of  Paris"  and,  on  treatment  with  water,  takes  up  its  water 
of  crystallization  once  more  and  "sets"  to  a  compact  mass  of  crystals. 

Calcium  or  magnesium  present  in  solution  as  sulfates  cannot  be  re- 
moved by  boiling.  Water  containing  these  elements  as  sulfates  is  said 
to  be  "permanently  hard,"  since  it  cannot  be  "softened"  except  by 
addition  of  chemical  reagents. 

Since  CaS04  is  so  much  more  soluble  than  CaC03,  on  boiling  with 
Na2C03,  the  former  is  readily  transposed  into  the  latter. 

CaS04  +  Na2C03  -»  CaC03  +  Na2S04 

The  chromate  ion,  Cr04,  produces  no  precipitate  except  in  concen- 
trated solutions,  for  CaCrO4  is  soluble  in  about  200  parts  of  H20.  Solu- 


STRONTIUM.  97 

tions  of  alkaline  dichromates  never  contain  sufficient  concentration  of 
Cr04  ions  to  cause  a  precipitate. 

The  chlorid  ion,  Cl,  produces  no  precipitate,  for  the  chlorid,  CaCL, 
like  the  bromid  and  iodid,  is  deliquescent  and  very  soluble.  The  anhy- 
drous salt  dissolves  in  about  1J  parts  of  H20. 

The  fluorid  ion,  F,  precipitates  white  CaF2,  soluble  in  about  27,000 
parts  of  water.  It  is  slowly  soluble  in  HC1  and  other  strong  acids,  for 
HF  is  a  weak  acid  (see  page  16).  It  occurs  in  nature  as  fluorspar  or 
fluorite. 

The  cyanid  ion,  ON,  causes  no  precipitate,  since  Ca(CN)2  is  readily 
soluble. 

The  phosphate  ion,  P04,  precipitates  white  Ca3(P04)2,  soluble  in 
about  40,000  parts  of  H20,  readily  soluble  even  in  acetic  acid.  By  acid 
sodium  phosphate,  HNa2P04,  which  is  the  ordinary  precipitating  re- 
agent, Ca3(P04)2  is  precipitated  only  in  ammoniacal  solutions.  In  neu- 
tral solutions,  acid  phosphate,  HCaP04,  falls.  Diacid  calcium  phos- 
phate, CaH4(P04)2,  is  the  " soluble  phosphate"  present  in  fertilizers. 

The  oxalate  ion,  C204,  precipitates  white  CaC204.2H20,  soluble  in 
about  140,000  parts  of  H20.  Since  oxalic  acid,  H2C204,  dissociates  to  a 
greater  degree  than  H(C2H302),  calcium  oxalate  is  insoluble  in  acetic 
acid,  although  its  precipitation  may  be  held  up  somewhat  if  the  concen- 
tration of  acetic  acid  becomes  too  great.  Calcium  oxalate  is  readily 
soluble  in  the  strong  mineral  acids,  however.  When  precipitated  from 
a  cold  solution  the  crystals  of  calcium  oxalate  are  so  minute  as  to  run 
through  a  filter  paper  readily  and  produce  a  milky  filtrate.  By  heating 
the  liquid,  either  previous  to  or  after  precipitation,  larger  crystals  are 
formed  which  may  be  filtered  without  difficulty. 

STRONTIUM.     Sr  — 87.6. 

Strontium  so  closely  resembles  calcium  in  all  essential  features  that 
a  detailed  description  is  unnecessary.  Its  nature  as  an  element  and  the 
properties  of  its  compounds  are  usually  intermediate  between  calcium 
on  the  one  side  and  barium  on  the  other. 

Strontium  salts  color  a  Bunsen  flame  an  intense  crimson.  As  the 
chlorid  is  more  volatile  and  easily  decomposed  than  other  salts,  the  best 
results  are  obtained  by  moistening  the  substance  with  HC1  before 
heating. 

With  the  spectroscope,  strontium  is  detected  by  the  presence  of  six 
bright  red  bands,  one  orange  and  one  blue  band.  (See  frontispiece.) 

The  only  strontium  ion  is  Sr.    For  reactions  of  this  ion,  solutions  of 

8 


98  QUALITATIVE    ANALYSIS. 

any  soluble  salt  will  serve.  The  nitrate,  Sr(N03)2.4H20,  is  efflorescent. 
The  anhydrous  salt  is  soluble  in  5  parts  of  H20  but  is  insoluble  in  boil- 
ing amyl  alcohol  (distinction  from  Ca). 

The  hydroxyl  ion,  OH,  produces  no  precipitate  except  in  very  con- 
centrated solutions,  for  Sr(OH)2.8H20  is  soluble  in  about  60  parts  of 
H20. 

The  carbonate  ion,  C03,  precipitates  white  SrC03,  soluble  in  about 
100,000' parts  of  H20.  It  is  similar  in  all  respects  to  CaC03  but  requires 
a  higher  temperature  to  decompose  it  into  CO2  and  SrO. 

The  sulfate  ion,  S04,  precipitates  white  SrSO4,  soluble  in  about  10,000 
parts  of  H20.  Since  CaS04  is  much  more  soluble  than  SrS04,  the  intro- 
duction of  a  saturated  solution  of  CaS04  into  solutions  containing  the 
Sr  ion  causes  a  precipitate  which  appears  somewhat  tardily,  inasmuch 
as  SrS04  does  not  separate  from  supersaturated  solutions  at  once.  This 
reaction  serves  to  distinguish  barium  and  calcium  from  strontium,  for 
the  former  gives  an  immediate  precipitate  and  the  latter  none  at  all. 

Since  SrS04  is  more  soluble  than  SrC03,  on  boiling  with  Na2C03, 
the  former  is  transposed  into  the  latter  compound. 

The  chromate  ion,  Cr04,  precipitates,  from  concentrated  solutions 
only,  yellow  SrCr04,  soluble  in  about  800  parts  of  H20.  Alkaline  di- 
chromate  solutions  contain  the  Cr04  ion  in  sufficient  concentration  to 
cause  a  precipitate  in  very  concentrated  solutions  of  strontium  salts. 

The  oxalate  ion,  C704,  precipitates  white  SrC204,  fairly  insoluble  in 
water  but  somewhat  soluble  in  acetic  acid. 

BARIUM.     Ba  — 137.4. 

The  metal  barium  closely  resembles  calcium  in  properties.  It  decom- 
poses water  readily  at  ordinary  temperatures,  forming  Ba(OH)2  and 
liberating  hydrogen. 

Barium  salts,  with  few  exceptions,  are  colorless.  In  all  of  its  com- 
pounds the  metal  is  bivalent. 

Barium  salts  color  a  Bunsen  flame  green.  Moistening  the  substance 
to  be  tested  with  HC1  aids  this  reaction  considerably,  owing  to  the  for- 
mation of  BaCl2  which  is  more  volatile  than  other  barium  salts.  Barium 
sulfate  is  not  appreciably  volatile  at  the  temperature  of  the  blowpipe 
and  is  not  decomposed  by  HC1,  therefore  gives  but  little  flame  coloration. 
Heated  on  a  platinum  wire  in  the  reducing  flame  it  is  converted  to  the 
sulfid  which,  on  treatment  with  HC1,  gives  a  satisfactory  indication 
when  reheated. 

With  the  spectroscope,  barium  is  indicated  by  a  number  of  bright 
green  lines.  (See  frontispiece.) 


BARIUM.  99 

The  only  barium  ion  known  is  Ba.  For  the  reactions  of  this  ion,  any 
soluble  salt  will  serve.  The  nitrate,  Ba(N03)2,  crystallizes  anhydrous 
from  solution.  It  is  soluble  in  about  11  parts  of  H2O  but  is  entirely  in- 
soluble in  boiling  amyl  alcohol  (like  Sr,  but  unlike  Ca).  The  acetate 
Ba(C2H302)2,  is  readily  soluble. 

The  hydroxyl  ion,  OH,  produces  no  precipitate,  for  Ba(OH)2.8H,O 
is  soluble  in  about  20  parts  of  cold  water  and  in  about  twice  its  own 
weight  of  boiling  water.  Barium  hydroxid  is  a  trifle  more  dissociated 
than  Ca(OH)2  at  corresponding  dilutions  and  its  solution  is  known  as 
"baryta  water." 

The  carbonate  ion,  C03,  precipitates  white,  amorphous  BaC03  which, 
on  heating,  becomes  crystalline.  It  is  soluble  in  about  50,000  parts  of 
H20.  In  the  presence  of  C02,  the  acid  carbonate  is  formed,  soluble  in 
about  1000  parts  of  H2O.  On  boiling  with  NH4C1,  the  carbonate  dis- 
solves and  the  chlorid  is  slowly  formed,  just  as  with  CaC03.  (See  page 
96.)  It  is  found  in  nature  as  the  mineral  witherite. 

Barium  carbonate  requires  a  much  higher  temperature — about  1500° 
—to  decompose  it  into  C02  and  the  oxid  than  does  CaC03,  being  more 
stable  under  the  influence  of  heat  than  any  other  carbonates  except  those 
of  the  fixed  alkalies.  Barium  oxid,  BaO,  like  CaO,  unites  with  water 
forming  the  hydroxid  which  readily  absorbs  C02  and  passes  into  car- 
bonate again. 

On  heating  in  a  current  of  air,  BaO  absorbs  oxygen  and  becomes 
Ba02.  At  a  higher  temperature  it' breaks  down  again  into  the  monoxid 
and  liberates  oxygen. 

Barium  peroxid,  treated  with  dilute  acids,  forms  hydrogen  peroxid, 
H202. 

The  sulfid  ion,  S,  causes  no  precipitate,  for  BaS  is  fairly  soluble  and 
hydrolyzes  to  soluble  products  as  does  CaS. 

The  sulfate  ion,  S04,  precipitates  white  BaS04,  soluble  in  about 
400,000  parts  of  water.  The  naturally-occurring  product  is  the  mineral 
baryte.  It  dissolves  in  concentrated  H2SO4  somewhat,  forming  the  acid 
sulfate,  BaH2(S04)2,  but  is  reprecipitated  on  diluting.  Hydrochloric 
acid  has  no  appreciable  solvent  effect.  Since  BaS04  is  so  much  more 
insoluble  than  BaCO3,  on  boiling  BaS04  with  Na2C03,  an  appreciable 
reaction  takes  place  only  when  the  concentration  of  Na2C03  approximates 
that  of  a  saturated  solution.  Even  then  the  metathesis  can  be  completed 
only  on  filtering  off  the  precipitate  and  treating  with  fresh  Na2CO3  re- 
peatedly, a  method  which  is  obviously  not  satisfactory.  By  fusion  with 


100  QUALITATIVE    ANALYSIS. 

4-5  times  its  bulk  of  Na2CO3,  BaSO4  may  be  more  readily  transposed, 
however. 

If  a  saturated  solution  of  CaS04  be  used  to  introduce  the  S04  ion 
into  solutions  containing  the  Ba  ion,  an  immediate  precipitate  results 
(distinction  from  Sr  and  Ca). 

The  chromate  ion,  Cr04,  precipitates  canary-colored  BaCr04,  soluble 
in  about  250,000  parts  of  H2O.  It  is  not  appreciably  soluble  in  acetic 
acid,  but  dissolves  readily  in  the  strong  acids.  Solutions  of  alkaline  di- 
chromates  contain  Cr04  ions  in  sufficient  concentration  to  cause  a  pre- 
cipitate of  BaCr04. 

The  chlorid  ion,  01,  produces  no  precipitate,  for  BaCl2.2H20,  like 
the  bromid  and  iodid,  is  very  soluble,  requiring  only  2  parts  of  water  for 
its  solution.  In  concentrated  HC1,  it  is  insoluble  and  if  sufficient  HC1 
be  added  to  a  solution  containing  Ba  ions,  a  precipitate  will  fall. 

The  fluorid  ion,  F,  precipitates  wrhite  BaF2,  which  is  somewhat  more 
soluble  in  water  than  CaF2. 

The  phosphate  ion,  P04,  precipitates  white  Ba3(P04)2,  soluble  even  in 
acetic  acid.  By  acid  sodium  phosphate,  HNa2P04,  the  ordinary  precipi- 
tating reagent  for  phosphates,  Ba3(P04)2  is  precipitated  only  in  alkaline 
solutions.  In  neutral  solutions  HBaP04  falls. 

The  oxalate  ion,  C204,  precipitates  white  BaC2O4,  soluble  in  about 
2500  parts  of  water.  In  the  strong  acids,  as  well  as  in  hot  acetic  acid, 
it  is  readily  soluble. 

The  silicofluorid  ion,  SiF6,  precipitates  white  BaSiF6,  soluble  in 
about  3500  parts  of  H2O  at  ordinary  temperatures.  In  dilute  HC1  it  is 
considerably  more  soluble,  however,  and  from  acid  solutions  requires  some 
standing  to  effect  appreciable  precipitation. 

Inasmuch  as  the  Ba  ion  is  an  important  reagent  in  detecting  and 
differentiating  between  acid  radicles  or  anions,  a  list  of  salts  formed  by 
the  Ba  ion  with  the  acids  considered  in  the  present  scheme  of  analysis  is 
appended,  together  with  the  solubilities  of  each  in  water  and  dilute  HCL 
(See  next  page.) 

Group  2,  Second  Type. 

Zinc  65.4 ;          Cadmium  112.4 ;          Mercury  200. 

The  members  of  this  group  are  heavy  metals,  each  possessing  a  bright 
metallic  luster.  They  do  not  pass  as  readily  into  the  ionic  condition 
as  the  first  type  metals  do ;  the  compounds  are  less  stable  and  from  them 
the  elements  may  be  obtained  without  difficulty.  Their  oxids  unite  with 
water  only  indirectly  to  form  hydroxids,  which  are  by  no  means  as 


BAKIUM    SALTS. 


101 


Acid. 

Formula. 

Color. 

H20. 

ECl. 

Acetic 

Ba(C2H302)2 

White 

Sol. 

Sol. 

Arsenic 

Ba3(AsO4)2 

White 

Insol. 

Sol. 

Arsenious 

Ba(AsO2)2 

White 

Insol. 

Sol. 

Boric 

Ba(B02)2 

White 

Insol. 

Sol. 

Carbonic 

BaCO3 

White 

Insol. 

Sol. 

Chloric 

Ba(C103)2 

White 

Sol. 

Sol. 

Chromic 

BaCrO4 

Yellow 

Insol. 

Sol. 

Cyanic 

Ba(OCN)2 

White 

Sol. 

Sol. 

Ferricyanic 

Ba3(FeC0N6)2 

Eed 

Sol. 

Sol. 

Ferrocyanic 

Ba2(FeC6N8) 

Yellow 

Sol. 

Sol. 

Hydrobromic 

BaBr2 

White 

Sol. 

Sol. 

Hydrochloric 

Bad, 

White 

Sol. 

Sol. 

Hydrocyanic 

Ba(CN)2 

White 

Sol. 

Sol. 

Hydrofluoric 

BaF2 

x  White 

Insol. 

Sol. 

Hydriodie 

BaI2 

White 

Sol. 

Sol. 

Hydrosilicofluoric 

BaSiF6 

White 

Insol. 

Insol. 

Hydrosulfuric 

BaS 

Lemon 

Sol. 

Sol. 

Hypochlorous 

Ba(ClO)a 

White 

Sol. 

Sol. 

Nitric 

Ba(N03)2 

White 

Sol. 

Sol. 

Nitrous 

Ba(NO2)2 

White 

Sol. 

Sol. 

Oxalic 

BaC2O4 

White 

Insol. 

Sol. 

Perchloric 

Ba(C104)2 

White 

Sol. 

Sol. 

Permanganic 

Ba(MnO4)2 

Purple 

Sol. 

Sol. 

Phosphoric 

Ba3(P04)2 

White 

Insol. 

Sol. 

Silicic 

BaSiOs 

White 

Insol. 

Insol. 

Sulfuric 

BaSO, 

White 

Insol. 

Insol. 

Sulfurous 

BaSO3 

White 

Insol. 

Sol. 

Thiocyanic 

Ba(SCN)2 

White 

Sol. 

Sol. 

Thiosulfuric 

BaS2O3 

White 

Sol. 

Sol. 

Tartaric 

Ba(C4H4O6) 

White 

Insol. 

Sol. 

strong  as  those  of  the  first  type.  They  act  as  weak  bases.  As  the  atomic 
weight  rises,  the  metals  become  more  resistant  to  the  atmosphere,  more 
readily  reduced  from  their  compounds  and  their  salts  tend  to  become 
colored.  The  group  is  generally  bivalent.  Mercury  forms  a  whole  series 
of  compounds  which  are  usually  considered  univalent,  however,  and 
cadmium  exhibits  some  slight  tendency  in  the  same  direction. 

ZINC.     Zn  —  65.4. 

Zinc  is  a  heavy,  bluish-white,  crystalline  metal.  It  tarnishes  but 
little  in  the  atmosphere,  owing  to  the  formation  of  a  compact  film  of  the 
oxid,  which  prevents  further  action.  When  highly  heated,  it  burns  with 
a  bluish- white  flame.  Commercial  zinc,  often  called  "spelter,"  contains 
small  amounts  of  carbon,  iron  and  copper.  "Zinc  dust,"  the  finely 
divided  metal  obtained  by  cooling  the  vapor,  always  contains  ZnO. 

Pure  zinc  dissolves  but  slowly  in  HC1  or  H2SO4.  The  commercial 
article  is  readily  soluble  in  acids,  however,  even  in  acetic  acid,  owing  to 
the  fact  that  the  impurities  form  with  the  metal  small  electrolytic  couples 
which  hasten  the  solution  of  the  positive  plate  (i.  e.,  the  zinc)  and  liberate 


102  QUALITATIVE    ANALYSIS. 

hydrogen  abundantly.  A  considerable  amount  of  energy  is  liberated 
during  the  passage  into  the  ionic  condition.  Consequently,  to  obtain 
the  metal  from  its  compounds  requires  the  expediture  of  a  large  amount 
of  energy  which  is  absorbed  during  the  reaction.  For  this  reason,  zinc 
salts  are  readily  reduced  only  in  the  *  *  dry  way  '  '  on  heating.  Zinc  is  not 
precipitated  from  solutions  of  its  salts  by  any  of  the  common  metals,* 
but,  on  the  contrary,  causes  most  other  metals  to  be  precipitated  from 
solutions  of  their  salts. 

Zinc,  like  aluminum  (see  page  115),  is  attacked  by  the  strong  alkalies, 
forming  zincates  according  to  the  reaction  : 

Zn  -f  2NaOH  ->  Na,Zn02  +  H  2 
or,  writing  the  ionic  equation: 
Zn      2 


Zinc  oxid,  ZnO,  like  other  oxids  generally,  is  formed  by  heating  the 
hydroxid,  carbonate,  nitrate,  sulfid,  acetate  and  salts  of  other  organic 
acids,  or  by  heating  other  salts  with  Na2C03.  It  is  a  white  substance 
turning  yellow  on  heating.  When  heated  before  the  blowpipe  after 
being  moistened  with  a  drop  of  dilute  Co(N03)2  a  green  color  appears, 
Rinnmann's  green,  which  is  probably  due  to  the  formation  of  cobalt 
zincate,  CoZnO2. 

Heated  before  the  blowpipe  with  Na2C03  on  charcoal,  all  zinc  com- 
pounds are  reduced  to  the  metal  which  is  immediately  volatilized  and 
oxidized,  forming  on  the  charcoal  a  coating  of  the  oxid,  yellow  when  hot 
and  white  when  cold. 

Zinc  salts  contain  the  metal  always  in  the  bivalent  condition.  It 
acts  as  a  base-forming  element  of  about  the  same  strength  as  ferrous 
iron,  cobalt  and  nickel.  Zinc  compounds  are  generally  colorless.  The 
salts  of  the  common  laboratory  acids  are  soluble  and  in  solution  are 
hydrolyzed  but  little;  hence  from  cold  solutions  zinc  hydroxid  is  not 
precipitated  by  the  addition  of  BaCO3  as  is  the  case  with  the  salts  of 
the  trivalent  metals  like  ferric  iron.  (See  page  225.) 

++ 

The  simple  zinc  ion  is  Zn.  Complex  ions,  zinc-ammonia  and  others, 
are  known,  also  the  zincate  ion,  Zn02,  in  which  zinc  acts  as  an  acid-form- 
ing element. 

For  the  reactions  of  the  Zn  ion,  any  soluble  simple  salt  will 
serve.  The  nitrate,  Zn(NO3)2.6H2O,  is  deliquescent  as  is  the  acetate, 

*  Under  proper  conditions  Mg  will  throw  Zn  out  of  solutions  of  its  salts. 


ZINC.  103 

Zn(C2H302)2.7H20.     Each  is  extremely  soluble  and  will  melt  in  its 

own  water  of  crystallization  at  temperatures  but  little  above  the  ordinary. 

The   hydroxyl   ion,    OH,    precipitates   white   Zn02H2   which   shows 

both  acidic  and  basic  properties.     (See  page  22.)     In  the  presence  of 

++  + 

acids,  it  separates  into  Zn  and  OH  ions,  the  latter  uniting  with  the  H 

ions  of  the  acids  to  form  water,  and  leaving  in  solution  the  ions 
necessary  to  form  zinc  salts.  On  the  other  hand,  the  presence  of  a 

high  concentration  of  OH  ions,  the  ionization  of  Zn02H2  into  OH  ions 

+  =  + 

is  repressed  and  H  and  Zn02  ions  are  formed  instead.     In  this  case  H 

ions  from  zinc  hydroxid  and  OH  ions  from  the  base  unite  to  form  H2O 
and  leave  in  solution  the  ions  necessary  to  form  zincates.  Alkali  zincates 
are  well-known,  readily  soluble  salts  of  the  type  indicated  by  the  formula 

K  — 0 

Zn 
K  — O 

Because  of  the  formation  of  these  salts,  all  zinc  compounds  except 
ZnS  dissolve  readily  in  the  strong  alkalies. 

Zinc  hydroxid  is  soluble  in  NH4OH  forming,  not  ammonium  zincate, 
but  complex  zinc-ammonia  compounds,  giving  complex  ions,  Zn(NH3)6, 
exactly  analogous  to  the  cobalt-ammonia  compounds  (see  page  229).  Be- 
cause of  the  formation  of  these  bodies,  all  zinc  compounds  except  ZnS  and 
Zn2(Fe  =  C6N6)  dissolve  readily  in  NH4OH. 

With  NH4C1  and  other  ammonium  salts,  zinc  hydroxid  and  other 
zinc  salts  form  double  salts  of  the  type,  (NH4)2ZnCl4,  which  in  fairly 
strong  solutions,  at  least,  dissociate  into  NH4  ions  and  ZnCl4.  On  dilut- 
ing, simple  ions  are  formed  by  decomposition  of  the  complex  ions.  These 
compounds  are  characteristic  of  ferrous  iron,  cobalt,  nickel,  manganese 
and  magnesium  as  well  as  of  zinc  and,  because  of  their  formation,  in 
the  presence  of  NH4C1  no  precipitate  is  given  by  any  of  these  metals* 
when  treated  with  NH4OH  or  (NH4)2CO3. 

The  carbonate  ion,  C03,  causes  a  precipitate  of  varying  composition, 
apparent aly  consisting  of  much  Zn(OH)2  with  some  ZnC03. 

The  sulfid  ion,  S,  precipitates  from  solutions  neutral,  alkaline,  or 
acid  with  acetic  acid  only,  white  ZnS,  readily  soluble  in  the  dilute  min- 
eral acids.  In  neutral  solutions,  the  precipitation  is  not  complete  be- 

*  Manganese  gives  a  precipitate  with  (NH4)2CO3  (see  page  217). 


104  QUALITATIVE    ANALYSIS. 

cause  of  the  liberation  of  mineral  acid  which  causes  the  reaction  to  come 
to  equilibrium,  as  indicated  by  the  equation: 

ZnCl2  +  H2S  **  ZnS  +  2HC1 

If  sodium  or  other  soluble  acetate  be  added,  the  precipitation  is 
complete.  By  the  interaction  of  the  strong  acid  with  the  acetate,  acetic 
acid  is  set  free.  In  acetic  acid  ZnS  is  insoluble. 

The  sulfate  ion,  S04,  causes  no  precipitate  since  the  sulfate  is  soluble 
in  its  own  weight  of  water.  From  solution  it  separates  with  7H20  in 
the  form  characteristic  of  the  "vitriols"  (see  page  188).  It  effloresces 
and  may  be  preserved  more  satisfactorily  than  other  zinc  salts  which  are 
generally  deliquescent.  As  "white  vitriol,"  it  finds  many  applications. 

The  chlorid  ion,  01,  produces  no  precipitate,  for  zinc  chlorid  is  ex- 
tremely soluble.  It  melts  readily  and  is  volatile  at  a  red  heat.  Its  solu- 
tions show  a  tendency  to  become  cloudy  because  of  the  formation  of  a 
basic  salt  due  to  hydrolysis. 

The  cyanid  ion,  ON,  precipitates  white  Zn(CN)2,  soluble  in  moder- 
ately concentrated  mineral  acids.  It  dissolves  readily  in  an  excess  of 
alkaline  cyanids,  forming  complex  cyanids  which  dissociate  into  alkali 
and  (ZnC4N4)  ions.  These  ions  are  quite  unstable  and  decompose  con- 
siderably into  the  simple  ions  in  solutions  of  ordinary  concentration. 

The  ferrocyanid  ion,  (Fe  =  C6N6),  precipitates  Zn2(Fe  =  C6N6),  white 
in  color  and  insoluble  in  dilute  acids  and  NH4OH.  By  alkalies  it  is  de- 
composed, however,  into  alkaline  zincates  and  ferrocyanids. 

The  ferricyanid  ion,  (Fe  =  C6N6),  precipitates  yellowish- white 
Zn3(Fe=C6N6)2,  insoluble  in  dilute  acids. 

The  phosphate  ion,  P04,  precipitates  white  Zn3(P04)2,  soluble  even 

in  acelic  acid. 

CADMIUM.     Cd  — 112.4. 

Cadmium  is  a  rather  soft,  white  metal,  resembling  zinc  very  closely 
in  appearance  and  properties.  It  is  not  appreciably  tarnished  by  the 
constituents  of  the  atmosphere  at  ordinary  temperatures  but,  when 
heated,  unites  with  oxygen,  producing  a  vigorous  combustion  and  giving 
off  clouds  of  brown  smoke  consisting  of  the  oxid,  CdO. 

The  metal  dissolves  slowly  in  dilute  mineral  acids,  HNO3  being,  as 
is  usual,  the  most  satisfactory  solvent.  In  the  passage  from  the  metallic 
into  the  ionic  condition  some  energy  is  liberated,  hence  cadmium  salts 
are  stable  and  not  particularly  easy  of  reduction.  They  are  generally 
colorless  and  for  the  most  part  insoluble.  Of  the  common  salts  the 
nitrate,  sulfate,  acetate  and  the  halids  are  readily  soluble. 


CADMIUM.  105 

Heated  before  the  blowpipe  with  Na2C03  on  charcoal,  all  cadmium 
salts  are  reduced  to  the  metal.  A  metallic  globule  is  not  obtained,  how- 
ever, because  cadmium  is  volatilized.  As  the  hot  vapor  moves  away  from 
the  heated  charcoal,  it  oxidizes  rapidly,  forming  a  brown  coating  on  any 
cold  surface. 

-j-4- 

The  ordinary  cadmium  ion  is  Cd,  furnished  in  abundance  by  solutions 
of  the  nitrate  or  other  soluble  salt  with  the  exception  of  the  halids.  These 
cadmium  compounds  dissociate  to  a  much  smaller  degree  than  do  salts 

generally  (see  footnote,  page  21).     Complex  ions  formed  by  the  union 

++  

of  the  simple  Cd  ion  with  NH3,  and  with  the  CN  ion  also,  are  well  known. 
The  hydroxyl  ion,  OH,  added  to  a  solution  containing  the  Cd  ion, 
produces  a  precipitate  of  white  Cd(OH)2.  It  is  insoluble  in  an  excess 
of  NaOH  or  other  reagent  furnishing  OH  ions  in  high  concentration, 
being  in  this  way  differentiated  from  two  other  white  hydroxids, 
Zn(OH)2  and  Pb(OH)2,  which  are  soluble.  In  the  presence  of  free  NH3, 

+  4- 

complex  ions  are  formed  of  which  the  best  known  is  Cd(NH3)4.  Since 
the  compounds  formed  by  this  ion  are  generally  soluble,  Cd(OH)2  dis- 
solves in  an  excess  of  NH4OH,  thus  showing  a  similarity  to  zinc  and 
copper  and  difference  from  lead. 

The  hydroxid  is  soluble  in  alkaline  cyanids  forming  soluble  complex 
cyanids. 

On  heating,  white  Cd(OH)2  breaks  down  into  brown  CdO  and  water, 
as  all  hydroxids  do. 

The  carbonate  ion,  C03,  precipitates  white  CdC03  which,  on  heating, 
breaks  down  into  the  brown  CdO,  as  all  except  carbonates  of  the  alkali 
metals  do. 

The  sulfid  ion,  S,  precipitates  yellow  CdS  from  neutral  and  alkaline 
solutions  and,  also,  from  solutions  that  contain  only  a  small  concentra- 

4-  + 

tion  of  H  ions.  The  sulfid  is  soluble  in  a  concentration  of  H  ions  greater 
than  is  given  by  a  4  per  cent,  solution  of  HC1,  hence  CdS  is  not  precipi- 
tated in  the  presence  of  much  strong  acid.  Cadmium  sulfid  is  not  solu- 
ble in  (NH4)2S.  xS  and  in  this  way  may  be  differentiated  from  the  yel- 
low sulfids  of  arsenic,  antimony  and  tin,  all  of  which  are  soluble  in  yellow 
ammonium  sulfid. 

From  slightly  acid  solutions  of  CdCl2,  an  orange  compound  having 
the  proportions  represented  by  the  formula  CdS.CdCl2  is  first  precipi- 
tated but  is  later  changed  into  the  pure  yellow  CdS. 

The  sulfate  ion,  S04,  produces  no  precipitate  since  CdS04  is  readily 
soluble.  The  crystallized  salt  ordinarily  has  the  formula  (CdS04)3. 
8H20  but  it  may  crystallize  in  the  characteristic  form  of  the  ' '  vitriols, ' ' 
CdS04.7H20. 


106  QUALITATIVE   ANALYSIS. 

The  chromate  ion,  Cr04,  causes  no  precipitate,  as  CdCr04  is  soluble. 

The  chlorid  ion,  01,  causes  no  precipitate,  as  CdCl2  is  soluble.  It 
crystallizes  with  2H20.  The  bromid  and  iodid  are  similar. 

The  cyanid  ion,  ON,  precipitates  white  Cd(CN)2,  soluble  in  acids, 
NH4OH  and  an  excess  of  alkaline  cyanid,  forming  a  double  cyanid  of 
the  type  Na2CdC4N4,  which  gives  rise  to  the  complex  ion,  (CdC4N4). 
This  ion  is  not  very  stable  and  even  in  concentrated  solutions  decom- 

-H- 

poses  to  a  considerable  extent  into  the  simple  Cd  and  CN.ions.  Hence, 
when  H2S  is  passed  into  the  solution,  CdS  is  precipitated,  as  is  always 

-H-  = 

the  case  when  Cd  and  S  ions  come  together  in  neutral  or  alkaline  solu- 
tions. In  this  particular,  cadmium  differs  from  copper  which  forms  a 
similar  double  cyanid  of  the  type  Na3CuC4N4,  giving  rise  to  the  complex 
cyanid  ion,  (CuC4N4).  The  latter  is  very  stable;  in  solutions  of  fair 
concentration  it  does  not  dissociate  appreciably  into  the  simple  ion  and, 
therefore,  is  not  precipitated  by  H2S. 

The  ferrocyanid  ion,  (Fe  =  C6N6),  precipitates  white  Cd2(Fe  = 
C6N6),  insoluble  in  dilute  acids.  Fairly  concentrated  acid  readily  dis- 
solves it  and  NH4OH  does  also. 

MERCURY   (Hydrargyrum).    Hg  — 200. 

Mercury  is  the  only  metallic  element  liquid  at  ordinary  tempera- 
tures. It  melts  at  —  39.4°  and  boils  at  357°.  It  is  very  heavy,  possesses 
a  silvery- white  lustre  and  alloys  readily  with  most  other  metals.*  Mer- 
cury alloys  are  known  as  amalgams. 

The  metal  does  not  tarnish  in  the  air  at  ordinary  temperatures  but 
does  unite  slowly  with  oxygen  when  near  the  boiling-point.  At  a  higher 
temperature  the  red  HgO  is  decomposed  into  the  elements. 

Mercury  is  but  little  soluble  in  HC1  or  dilute  H2S04  for  two  reasons. 
First,  the  metal  does  not  pass  easily  into  the  ionic  condition  and,  there- 
fore, does  not  take  the  charge  from  hydrogen  ions  readily.  Secondly, 
what  little  action  HC1  and  H2SO4  have  on  the  element  results  in  the 
formation  of  insoluble  coatings  of  (HgCl)2  or  Hg2SO4,  which  protect 
the  metal  from  further  action.  Oxidizing  agents,  such  as  HN03  and 
hot  concentrated  H2S04,  dissolve  the  metal  readily,  however.  (Note 
similar  action  of  acids  on  copper,  page  79.) 

When  mercury  and  HNO3  are  brought  together,  the  temperature  and 
the  relative  proportions  of  acid  and  metal  determine  the  reaction  prod- 
uct. If  mercury  is  present  in  excess  and  the  acid  is  dilute  and  cold, 
mercurous  nitrate,  (HgN03)2,  is  formed.  If  the  HNO3  is  in  excess, 

*  Iron  and  platinum  do  not  alloy  with  mercury  at  ordinary  temperatures. 


MERCUEY    COMPOUNDS.  107 

(HgN03)2  is  oxidized  to  mercuric  nitrate,  Hg(NO3)2,  especially  if  the 
acid  is  at  all  concentrated  or  heated. 

Since  mercury  shows  a  disinclination  to  pass  into  the  ionic  condition 
and  absorbs  considerable  energy  when  it  assumes  the  electric  charge, 
many  other  metals  will  precipitate  the  element  from  solutions  of  its 
salts  because  these  metals  show  a  stronger  tendency  to  form  ions  than 
mercury  does.  Zinc,  iron  and  copper  precipitate  the  metal  with  great 
readiness.  Mercury  will  throw  down  metallic  silver,  gold  and  platinum, 
however,  from*  solutions  of  their  salts. 

Mercury  salts,  for  this  reason,  are  easily  reduced  to  the  metal  by 
strong  reducing  agents.  Heat  alone  is  sufficient  to  decompose  them  i» 
most  instances.  On  the  contrary,  quite  a  number  of  mercury  salts  sub- 
lime without  decomposition.  By  the  addition  of  Na2CO3,  however,  all 
salts  decompose  when  heated  in  a  closed-tube,  and  mercury  distills  out 
of  the  mixture  and  collects  on  the  cooler  portions  of  the  tube. 

Mercury  salts  are  of  two  general  types;  mercuric,  in  which  the  ele- 
ment is  bivalent,  and  mercurous,  in  which  the  element  may  be  considered 
as  univalent.  At  ordinary  temperatures,  however,  the  molecules  of  mer- 
curous salts  seem  to  be  doubled.  Mercurous  nitrate,  for  instance,  is  rep- 
resented by  the  formula  (HgN03)2  rather  than  by  HgN03,  and  other 
salts  have  corresponding  double  molecules.  (See  also  page  83.) 

Mercuric  compounds  are  stable  and  frequently  colorless.  Salts  of 
the  common  laboratory  acids  are  soluble  and  generally  separate  from 
solution  without  water  of  crystallization. 

Mercurous  compounds  are  often  unstable  and  show  a  decided  ten- 
dency to  change  into  the  mercuric  condition.  They  are  usually  much 
less  soluble  than  the  corresponding  mercuric  forms. 

Mercury  ions  are  of  two  types  corresponding  to  the  two  classes  of 

-f  -H- 

salts,  the  mercurous  ion,  (Hg)2,  and  the  mercuric  ion,  Hg,  being  the 
most  important.  Since  the  compounds  formed  by  the  mercurous  ion 
are  entirely  distinct  from  the  compounds  of  the  mercuric  ion,  they  will 
be  treated  separately. 

The  Mercuric  Ion,  Hg. 

For  the  reactions  of  the  mercuric  ion,  solutions  of  the  sulfate,  acetate 
or  chlorid  may  be  used,  but  the  nitrate  is  possibly  most  satisfactory.  It 
crystallizes  from  solutions  in  the  form  (Hg(N03)2)2.H20. 

The  hydroxyl  ion,  OH,  added  to  a  solution  containing  the  Hg  ion, 
doubtless  forms  Hg(OH)2  in  the  first  place,  but  this  compound  breaks 
down  immediately  into  H20  and  yellow  HgO.  This  precipitated  prod- 


108  QUALITATIVE    ANALYSIS. 

uct  differs  in  color  from  a  red  compound  having  the  same  composition 
formed  by  heating  the  metal  in  air.  By  boiling  with  an  excess  of  strong 
alkali,  the  yellow  modification  is  changed  over  into  the  red.-  Mercuric 
oxid  turns  black  on  heating,  but  becomes  red  again  on  cooling. 

If  NH4OH  is  used  as  a  precipitant,  a  white  precipitate  falls  in  place 

44. 

of  the  yellow  HgO,  for  the  Hg  ion  reacts  with  the  NH3  (which  is  always 
present  as  a  decomposition  product  of  NH4OH)  to  form  complex  ions 
of  which  the  mercuric-amido,  HgNH2  is  the  best  known.  Its  formation 
takes  place  according  to  the  equation  : 

H 

Cl  H  N—  H 

Hg         +H  —  N        ^Hg  +HC1 

\  \  \ 

Cl  H  Cl 

or,  writing  the  ionic  equation  : 


Since  the  salts  formed  by  this  ion  are  more  insoluble  than  HgO,  they  are 
formed  in  preference.  Mercuric  amido-chlorid,  HgNH2Cl,  taken  as  a 
typical  salt,  is  white  and  insoluble  in  acids. 

The  carbonate  ion,  C03,  precipitates  yellow-brown  HgC03,  an  un- 
stable compound  which,  slowly  at  ordinary  temperatures,  rapidly  on 

heating,  loses  its  C02  and  becomes  HgO.     If  (NH4)2C03  be  used  as  a 

++ 
precipitant,  the  Hg  ion  reacts  by  preference  with  the  NH3  present  in 

solution  because  of  hydrolysis  (see  page  76)  and  the  same  salt  is  pre- 
cipitated as  with  NH4OH. 

The  sulfid  ion,  S,  precipitates  black  HgS  from  alkaline  and  acid  solu- 
tions as  well.  In  the  presence  of  considerable  acid,  especially  HC1,  a 
white  precipitate  having  the  proportionate  composition  HgCl2.2HgS  first 
falls.  On  further  treatment  with  H2S,  this  passes  through  yellow, 
orange  and  red  to  the  black  sulfid. 

On  heating,  HgS  sublimes  and  changes  to  the  red  form,  cinnabar,  the 
sulfid  found  in  nature  and  known  as  the  pigment  "vermilion." 

Mercuric  sulfid  is  insoluble  in  acids,  alkalies  and  ammonium  sulfid. 
It  is  readily  soluble  in  the  sulfids  of  the  "fixed"  alkali  metals,  however, 
according  to  the  equation  : 

S-»K2HgS2 


MERCURIC    SALTS.  109 

In  concentrated  solution  the  ions  of  this  double  sulfid  are  K  and  HgS2. 
On  diluting,  HgS2  decomposes  into  the  simple  ions  Hg  and  S  and  HgS 
is  again  precipitated. 

By  long  continued  action  of  boiling  HN03  on  HgS,  it  may  be  par- 
tially oxidized  and  form  white  Hg(N03)2.2HgS,  Aqua  regia  oxidizes 
this  white  compound  as  well  as  the  black  HgS  readily  to  HgCl2  and 
H2S04. 

The  sulfate  ion,  S04,  causes  no  precipitate,  since  HgS04  is  readily 
soluble.  With  all  the  mercury  salts  of  acids  containing  oxygen,  there 
is  a  tendency,  more  marked  in  the  mercurous  than  in  the  mercuric  con- 
dition, toward  the  formation  of  insoluble  oxysalts  of  the  type  HgO. 
HgS04.  In  dilute  solutions  free  from  acid,  the  mass  action  of  the  water 
causes  appreciable  hydrolysis  of  the  salts  into  the  weak  base,  Hg(OH)2, 
which  decomposes  into  HgO,  as  previously  discussed.  This  combines 
with  the  normal  salt  to  form  insoluble  oxysalts,  which  are  thus  seen  to 
be  dehydrated  basic  salts  (see  page  21).  The  addition  of  a  little  acid 
to  the  solution  prevents  the  mass  action  of  the  water,  however. 

The  chromate  ion,  Cr04,  precipitates  yellow  HgCr04  from  concen- 
trated solutions,  or  possibly  a  red  oxysalt.  Both  dissolve  in  the  strong 
acids. 

The  chlorid  ion,  01,  causes  no  precipitate  since  HgCl2,  "corrosive 
sublimate,"  is  soluble  in  about  15  parts  of  water  at  ordinary  tempera- 
tures and  in  less  than  two  parts  at  100°. 

The  bromid  ion,  Br,  gives  no  precipitate  since  HgBr2  is  soluble. 

The  iodid  ion,  I,  precipitates  yellow  HgI2,  changing  rapidly  to  a 
red  modification.  Both  forms  are  soluble  in  an  excess  of  an  alkaline 
iodid  forming  a  complex  ion,  the  alkaline  salts  of  which  are  soluble.  A 
typical  salt  of  the  proportionate  composition  HgI2.2KI  dissociates  into 
K  and  HgI4  ions,  showing  that  a  new  compound  has  actually  been 
formed.  An  alkaline  solution  of  this  salt  turns  brown  in  the  presence 
of  traces  of  NH3  and,  as  the  "Nessler  reagent,"  is  used  to  detect  small 
amounts  of  ammonia  in  drinking  water.  All  the  mercury  halids  show 
a  great  tendency  toward  the  formation  of  complex  ions  of  this  same  type 
with  the  halids  of  other  metallic  elements.  (See  also  similar  lead  salts, 
page  145.) 

The  cyanid  ion,  CN,  precipitates  white  Hg(CN)2  from  concentrated 
solutions  containing  the  Hg  ion,  provided  that  halogen  ions  be  not 
present.  If  KCN  be  added  to  HgCl2,  no  precipitate  falls,  owing  to  the 
formation  of  some  one  of  a  series  of  double  salts  similar  to  those  noted 
above. 


110  QUALITATIVE    ANALYSIS. 

The  mercuric  halids  are  of  interest  because  of  the  fact  that  they 
ionize  to  a  very  small  degree,  a  very  marked  point  of  difference  between 
the  behavior  of  these  and  of  most  other  salts.  Mercuric  cyanid  separates 
so  few  ions  that  a  saturated  solution  conducts  the  electric  current  only  a 
trifle  better  than  pure  water.  The  tendency  toward  dissociation  is  some- 
what greater  in  the  iodid  and  increases  as  the  atomic  weight  of  the 
halogen  decreases.  (See  footnote,  page  20.) 

The  stannous  ion,  Sn,  in  SnCl2  solution,  causes  the  precipitation 
of  white  (HgCl)2  or  black  Hg,  depending  on  the  proportions  in  which 
the  reagents  are  present.  The  changes  are  caused  by  reduction,  as  the 
equations  show: 

2HgCl2  +  SnCl2  ->  (HgCl)2  +  SnCl4          HgCl2  +  SnCl2  -^  Hg  +  SnCl, 
2Hg  +  Sn  +  201  ->  (HgCl)2  +  Sn  Hg  +  Sn -> Hg  +  Sn" 

+  ^  4-  X  +*-  "^ 

+  u  The  Mercurous  Ion,  (Hg)2. 

The  mercurous  ion,.  (Hg)2,  is  found  in  appreciable  quantities  only  in 
solutions  of  the  nitrate,  (HgN03)2,  which  is  the  only  readily  soluble 
compound  formed  by  this  ion.  It  separates  from  solutions  with  two 
molecules  of  water  of  crystallization.  Because  of  hydrolysis,  neutral 
solutions  on  standing  separate,  as  a  yellow  powder,  a  basic  nitrate, 
( HgOH )  2  ( HgN03 )  2.  The  formation  of  this  precipitate  may  be  prevented 
by  the  addition  of  a  little  HN03.  In  order  to  prevent  oxidation  to  the 
mercuric  condition,  metallic  mercury  must  be  present  also,  the  excess 
of  metal  maintaining  the  salt  in  the  lower  form. 

The  hydroxyl  ion,  OH,  added  to  a  solution  containing  the  mercurous 
ion,  (Hg)2,  probably  forms  in  the  first  place  (HgOH)2,  but  this  immedi- 
ately separates  water  and  precipitates  black  Hg2O.  Mercurous  oxid  is 

very  unstable,  however,  and  soon  breaks  down  into  HgO  and  black  Hg. 

+  ++ 

With  NH4OH,  the  (Hg)2  ion  reacts  analogously  to  the  Hg  ion.     A 

mercurous  amido-chlorid,  or  other  amido-salt,  is  first  formed  which  is 
black  in  color.  Like  most  mercurous  salts,  it  decomposes  into  the  cor- 
responding mercuric  compound  and  the  free  metal,  as  the  equations 
indicate : 

H 

N 

/   \ 
/  H 

-*    Hg  +    Hg 

Cl 


MERCUROUS    SALTS.  Ill 

The  black  precipitate  is  insoluble  in  acids,  but  is  oxidized  by  aqua  regia 
to  mercuric  chlorid  according  to  the  equation : 

H 

Hg— N         +  10C1  -»  4HgCl2  +  N2  +  4HC1 


2 


L 


The  carbonate  ion,  C03,  precipitates  yellow  Hg2C03  which  soon  dark- 
ens, owing  to  decomposition  into  Hg2O  and  CO2  and  then  further  into 
HgO  and  Hg.  Ammonium  carbonate  causes  the  same  precipitate  as 
ammonium  hydroxid  because  of  the  NH4OH  formed  by  the  hydrolysis 
of  the  carbonate. 

The  sulfid  ion,  S,  precipitates  black  Hg2S.  Since  mercurous  sulfid  is 
not  stable  above  zero,  it  ordinarily  decomposes  into  HgS  and  Brg. 

The  sulfate  ion,  SO4,  precipitates  white  Hg2SO4  which,  being  some- 
what soluble,  may  not  fall  in  dilute  solutions. 

The  chromate  ion,  CF04,  causes  red  Hg2Cr04  to  fall,  soluble  in  HN03 
but  insoluble  in  alkaline  hydroxids. 

The  chlorid  ion,  Cl,  causes  a  granular  precipitate  of  white  (HgCl)2, 
soluble  in  about  300,000  parts  of  H2O  and  only  a  little  more  soluble  in 
acids  and  alkqj^j.  Mercurous  chlorid  is  known  medicinally  as  calomel. 
It  is  changed  i^  NH4OH  into  the  black  amido-chlorid  noted  above. 
Aqua  regia  oxidizes  (HgCl)2  into  the  soluble  HgCl2. 

The  bromid  ion^Br,  precipitates  yellowish-white  (HgBr)2. 

The  iodid  ion,  I,  causes  green  (Hgl)2  to  precipitate.  The  original 
color  soon  darkens  considerably,  owing  to  the  separation  of  metallic  mer- 
cury and  the  formation  of  HgI2,  which  dissolves  in  an  excess  of  an  alka- 
line iodid. 

The  cyanid  ion,  ON,  forms  in  the  first  place  (HgCN)2,  but  this  imme- 
diately decomposes  into  Hg(CN)2  and  Hg. 

The  stannous  ion,  Sn,  from  SnCl2  solution,  at  first  causes  the  precipi- 
tation of  (HgCl)2  because  of  the  presence  of  the  Cl  ion.  Subsequently 
the  Sn  ion  reduces  white  (HgCl)2  to  black  Hg.  It  is  itself  oxidized  to 

+-H-+ 

the  Sn  ion  at  the  same  time.     (See  page  110.) 

Group  3,  First  Type. 
Boron  11;          Aluminum  27.1;          Scandium  44.1;          Yttrium  89; 

Lanthanum  138.9. 

This  group  consists  of  trivalent  elements,  in  nature  intermediate 
between  the  strong  base-formers  and  the  strong  acid-formers.  The  hy- 


112  QUALITATIVE    ANALYSIS. 

drids  are  unimportant.  The  oxids  of  the  elements  with  low  atomic 
weights  are  anhydrids  of  weak  acids,  while  the  oxids  of  the  elements 
with  higher  atomic  weights  unite  with  water,  only  indirectly,  to  form 
hydroxids  which  are  weak  bases.  Compounds  are  generally  colorless 
and  stable,  and  it  is  only  with  difficulty  that  the  element  may  be  ob- 
tained from  them. 

The  first  two  elements  only  are  at  all  common.  The  hydroxid  of 
boron  acts  normally  as  an  acid.  Aluminum  hydroxid  acts  characteris- 
tically as  a  weak  base,  although  it  sometimes  shows  acid  properties.  This 
striking  difference  only  illustrates  the  general  rule  to  be  observed  in  all 
groups,  however,  viz.,  that,  as  the  atomic  weight  rises,  base-forming  prop- 
erties develop  and  acid-forming  qualities  disappear. 

BORON,  B  — 11. 

The  element  boron  is  known  as  an  amorphous  brown  powder  which 
does  not  melt  at  the  highest  temperature  of  the  electric  arc.  It  dissolves 
considerably  in  molten  aluminum,  from  which  it  separates  in  the  form 
of  exceeding  hard  and  brilliant  crystals  that  are  much  less  reactive  than 
the  amorphous  modification.  At  ordinary  temperatures  either  form  is 
stable  but,  when  heated,  each  burns  to  the  white  oxid,  B203. 

Boron  is  not  appreciably  attacked  by  HC1.  Hot  concentrated 
H2S04  or  HN03  react  with  it,  forming,  possibly,  a  nitrate  or  sulfate  (see 
also  page  156),  but  these  compounds  (if  they  exist)  are  completely  hydro- 
lyzed  into  B(OH)3  and  the  respective  acids  on  diluting  the  solution. 

Boron  acts  as  a  trivalent,  acid-forming  element,  much  weaker  than 
nitrogen,  phosphorous,  arsenic  or  carbon.  Its  acid  properties  are  more 
strongly  developed  than  those  of  aluminum,  however.  The  hydroxid, 
B(OH)3,  shows  practically  no  basic  properties  and,  therefore,  is  known 
as  boric  acid.  All  compounds  of  boron  of  any  importance  are  deriva- 
tives of  boric  acid. 

Compounds  of  boron  give  a  distinct  green  color  to  a  flame,  provided 
the  temperature  be  sufficiently  hot  to  volatilize  them.  A  Bun  sen  flame 
is  scarcely  hot  enough  and  a  blowpipe  is  needed  to  develop  this  color. 
On  moistening  with  H2S04,  most  boron  compounds  will  be  sufficiently 
decomposed  into  boric  acid  to  give  this  reaction,  but  some  silicates  con- 
taining boron  require  the  stronger  decomposing  action  of  hydrofluoric 
acid.  (See  page  135.) 

Boric  Acid. 

Ortho-boric  acid,  H3B03,  has  such  weakly  acidic  properties  that  its 
solution  conducts  the  electric  current  but  little  better  than  distilled  water 


BOEATES.  113 

and  gives  only  faint  reaction  to  litmus.  Its  solubility  increases  rapidly 
with  the  temperature  and  from  boiling  aqueous  solutions,  on  cooling,  it 
separates  in  gleaming  white  scales.  On  heating  to  100°,  H3B03  loses 
water,  becoming  meta-boric  acid,  HB02.  At  160°,  from  four  molecules 
of  HB02  one  molecule  of  water  passes  off,  forming  H2B407,  tetra-boric 
acid,  which,  at  a  higher  temperature,  breaks  down  into  the  anhydrid 
B203,  a  deliquescent  vitreous  substance. 

On  standing  with  water,  all  of  these  dehydration  products  become 
hydrated  again,  so  that  ortho-boric  acid  is  the  form  ordinarily  existing 
in  and  obtained  from  solution. 

Salts  of  ortho-boric  acid  are  not  definitely  known.  If  they  exist  at 
all,  they  are  decomposed  by  water  into  free  boric  acid  and  salts  of  meta- 
and  tetra-boric  acids.  The  conditions  of  the  experiment  determine  the 
character  of  the  precipitate  that  will  fall  when  a  salt  of  a  metal  is  added 
to  a  solution  containing  boric  acid  ions.  From  dilute  solutions  salts  of 
the  meta-  type  are  precipitated  most  generally,  independent  of  the  nature 
of  the  boric  acid  compound  used  in  making  the  solution,  thus : 

Na2B407  +  CaCl2  +  3H20  -»  Ca(B02)2  +  2NaCl  +  2H3B03 

Sodium  tetra-borate,  or  borax,  Na2B407.10H20,  is  the  most  important 
salt  of  boric  acid.  On  heating,  it  intumesces  as  the  water  boils  out  and 
then  fuses  to  a  clear,  colorless  glass,  known  as  "  borax  glass."  This 
anhydrous  borax  has  the  power  of  dissolving  metallic  oxids,  probably 
forming  metaborates,  many  of  which  are  very  characteristically  colored. 
A  typical  reaction  might  be  written : 

Na2B407  +  CoO  ->  2NaB02  +  Co(B02)2 

Borax-bead  reactions  furnish  important  information  in  blowpipe  analy- 
sis. (See  page  247.) 

Borates  of  the  alkali  metals  are  soluble  in  water.  Other  borates  are 
but  little  soluble  in  water ;  but,  being  salts  of  a  very  weak  acid,  they  are 
readily  dissolved  by  dilute  inorganic  acids.  Ammonium  chlorid  has  a 
decided  solvent  action,  exerting  a  strong  influence  in  holding  up  precipi- 
tates, often  preventing  their  formation  even  in  ammoniacal  solutions. 

Borates  are  readily  transposed  by  boiling  with  sodium  carbonate; 
barium  meta-borate,  for  instance,  giving  barium  carbonate  and  sodium 
borate. 

Salts  of  boric  acid,  as  well  as  the  anhydrid,  are  extremely  stable,  and 
are  not  volatile  except  at  high  temperatures.  Consequently,  on  fusion, 
neither  the  acid  nor  its  salts  are  removed  or  appreciably  changed.  Curi- 

9 


114  QUALITATIVE    ANALYSIS. 

ously  enough,  however,  boric  acid  is  volatile  in  steam  and  significant  loss 
may  occur  by  the  prolonged  boiling  of  an  acid  solution  before  it  comes 
to  dryness. 

Boric  acid  unites  with  the  organic  base,  methyl  alcohol,  forming  a 
salt,  methyl  borate,  which  is  readily  volatile. 


CH3— 0— H  H— 0  CH3— 0 

\  \ 

CH3— 0— H     +     H— 0— B     -»    CH3— 0— B     +     3H20 

/  / 

CH3— 0— H  H— 0  CH3— 0 

This  reaction  takes  place  most  satisfactorily  in  the  presence  of  con- 
centrated sulfuric  acid  which  not  only  sets  free  the  boric  acid,  but  also 
unites  with  the  water  formed  during  the  process,  thus  preventing  a 
reversal  of  the  reaction.  On  warming,  methyl  borate  passes  off  with 
any  excess  of  alcohol.  On  burning  the  vapor,  methyl  borate  is  oxidized 
to  boric  acid,  carbon  dioxid  and  water,  thus: 

2  ( CH3)3B03  +  902  ->  B203  +  9#20  +  6C02 

The  boric  anhydrid  thus  formed  gives  the  characteristic  green  boron 
flame.  Since,  however,  the  temperature  necessary  for  the  development 
of  this  color  is  quite  high,  it  is  more  satisfactory  to  burn  the  vapor  in  an 
apparatus  which  is  essentially  an  improvised  Bunsen  burner.  Into  the 
mouth  of  the  test-tube  containing  a  borate,  methyl  alcohol  and  sulfuric 
acid,  fit  a  stopper  pierced  by  a  glass  tube  which  is  drawn  to  a  fine  point. 
Support  above  this  orifice  a  piece  of  glass  tubing  two  to  four  inches  long 
and  a  quarter  of  an  inch  in  diameter,  in  such  a  position  that  when  the 
vapors  issue  from  the  test-tube  in  the  form  of  a  jet,  they  will  draw  into 
the  upper  tube  a  supply  of  air.  The  mixture  may  be  inflamed  at  the 
top  of  the  upper  tube,  the  increased  amount  of  air  producing  a  more 
perfect  combustion,  and  developing  a  higher  temperature,  as  in  the  case 
of  the  ordinary  Bunsen  burner. 

When  turmeric  paper  is  dipped  into  a  solution  containing  a  borate 
acidified  with  HC1  and  is  then  dried  at  the  temperature  of  boiling  water, 
the  yellow  color  of  the  turmeric  paper  is  changed  to  pink  or  red.  Alka- 
line solutions  change  the  red  to  a  greenish-black  color. 

+-I- 

The  barium  ion,  Ba,  precipitates  from  solutions  containing  the  borate 
ion,  white,  floccy  barium  meta-borate,  Ba(B02)2.  It  is  readily  soluble 
in  acids  and  in  ammonium  chlorid. 

+ 

The  silver  ion,  Ag,  precipitates  white  silver  meta-borate,  Ag(B02)2, 
soluble  in  dilute  nitric  acid  and  and  in  ammonium  hydroxid. 


ALUMINUM.  115 

ALUMINUM.    Al  — 27.1. 

Aluminum  is  a  light,  silver-white,  malleable  metal.  It  is  apparently 
not  acted  on  by  the  atmosphere  but,  in  reality,  a  thin  compact  coating 
of  aluminum  oxid  is  readily  formed  which,  being  very  resistant  to  chem- 
ical agents  generally,  protects  the  metal  from  further  action. 

The  metal  dissolves  readily  in  HC1  and  slowly  in  H2SO4.  In  HNO3 
it  is  not  appreciably  soluble,  however.  This  is  possibly  due  to  the  oxi- 
dizing action  of  HN03,  resulting  in  the  formation  of  a  protective  coating 
of  A1203.  In  solutions  of  the  strong  basic  hydroxids,  the  metal  dissolves 
with  the  liberation  of  hydrogen  and  formation  of  alkaline  aluminates 
(See  also  page  102.) 

2A1  +  6NaOH  -»  2Na3A103  +  3H2 

With  the  exception  of  magnesium  and  the  alkali  metals,  no  element 
shows  a  greater  tendency  to  pass  into  the  ionic  condition.  Under  appro- 
priate conditions,  practically  all  the  other  metals  are  precipitated  from 
solutions  of  their  salts.  On  the  other  hand,  no  metal  or  other  reducing 
agent  is  able  to  precipitate  aluminum  from  solutions  of  its  salts.  Even 
in  the  ' '  dry  way, ' '  the  reduction  is  accomplished  only  at  very  high  tem- 
perature and  with  considerable  difficulty. 

Aluminum  oxid,  * '  alumina, ' '  A1203,  is  a  firm,  white  substance  which, 
after  strong  heating,  becomes  very  resistant  to  acids  and  requires  fusion 
with  an  alkaline  carbonate  or  acid  sulfate  to  get  it  into  solution.  In 
nature,  it  occurs  as  the  minerals,  corundum  or  emery,  and  when  clear  and 
colored  slightly  with  other  oxids,  as  the  gems,  sapphire  and  ruby. 

Aluminum  compounds  are  generally  colorless.  The  salts  of  the 
common  laboratory  acids  are  all  soluble  and  they  hydrolyze  decidedly 
in  solution,  since  aluminum  is  a  very  weak  base-forming  element.  To  the 
fact  that  chromium  and  iron  exist  in  the  trivalent  condition  also  and  act  as 
very  weak  base-forming  elements  is  to  be  attributed  largely  the  similar 
solubility  relations  because  of  which  these  three  metals  are  classed  to- 
gether in  qualitative  analysis.  A  comprehensive  survey  of  the  proper- 
ties of  each  element  will  not  demonstrate  any  close  relationship  be- 
tween aluminum  and  chromium  and  iron. 

Heating  before  the  blowpipe  on  charcoal  with  Na2C03  produces 
only  the  oxid.  If  this  be  moistened  with  a  drop  of  dilute  Co(N03)  and 
reheated,  a  bright  blue  color,  Thenard's  blue,  appears.  Cobalt  nitrate, 
like  all  nitrates,  is  decomposed  on  heating  and  forms  the  oxid  which 
unites  with  A1203  to  form  the  blue  compound  which  is  probably  cobalt 
meta-aluminate,  Co(A102)2.  With  the  exception  of  the  phosphate  and 
silicate  which  require  Na2C03  to  decompose  them,  all  aluminum  salts 


116  QUALITATIVE    ANALYSIS. 

break  down  into  A1203  on  ignition.  Except  with  the  phosphate  and 
silicate,  therefore,  this  test  may  be  made  more  satisfactorily  by  heating 
on  porcelain  or  platinum  without  Na2C03. 

Of  aluminum  ions  the  only  kathion  is  Al.  The  metal  enters  as  a  con- 
stituent in  certain  anions  of  which  the  ortho-aluminate  ion,  AlO3,  and  the 
meta-aluminate  ion,  A102,  are  the  most  important. 

A  solution  of  any  of  the  ordinary  soluble  salts  of  aluminum  will  serve 
satisfactorily  for  the  study  of  the  reaction  of  the  Al  ion.  The  double 
sulfate  of  potassium  and  aluminum,  ordinary  alum,  may  also  be  used. 

The  hydroxyl  ion,  OH,  precipitates  white,  gelatinous  A1(OH)3,  which 
shows  considerable  tendency  to  form  colloidal  solutions.  This  is  readily 
prevented  by  the  presence  of  ammonium  salts,  however.  When  first 
precipitated,  the  hydroxid  is  readily  soluble  in  dilute  acids  but,  on  stand- 
ing in  the  solution,  becomes  less  and  less  soluble,  probably  due  to  partial 
dehydration,  for  the  oxid  is  much  less  soluble  than  the  hydroxid.  The 
properties  of  the  naturally-occurring  hydroxids  confirm  this  supposition. 
Hydrargillite,  A1(OH)3,  is  readily  soluble  in  hot  HC1  but  beauxite, 
A120(OH)4,  and  diaspore,  AIO(OH),  are  insoluble  in  the  same  solvent. 

Aluminum  hydroxid  possesses  weak  acid  as  well  as  basic  properties 
(see  page  22),  and  will  dissolve  in  solutions  of  the  strong  alkalies,  form- 
ing aluminates  of  which  two  types  are  recognized. 

Na— O 

\ 
Na— O— Al  Na— 0— Al  =  O 

/ 

Na— O 

Sodium  ortho-aluminate.  Sodium  meta-aluminate. 

Ammonium  aluminate  apparently  exists,  for  A1(OH)3  will  dissolve 
somewhat  in  NH4OH.  The  ammonium  salt  hydrolyzes  very  greatly  and 
precipitates  A1(OH)3,  especially  on  warming.  This  action  is  favored 
by  the  fact  that  NH4OH  decomposes  under  these  conditions  into  NH3  and 
H,0  and  may  be  boiled  off.  This  reduces  the  amount  of  NH4OH  below 
the  concentration  necessary  to  maintain  equilibrium,  hence  the  reaction 
ceases  to  be  reversible  and  all  of  the  aluminum  may  be  precipitated  from 
solution. 

(NH4)3A1O3  +  3H20  ->  A1(OH)3  +  3NH,  +  3H20 

The  analogous  precipitation  of  A1(OH)3  from  solutions  of  sodium 
or  potassium  aluminate  can  not  be  accomplished,  for  the  A1(OH)3  and 
alkali,  as  fast  as  formed,  interact  to  regenerate  the  aluminate  again. 

Na3 A1O3  +  3H2O  *»  Al  ( OH )  3  +  3NaOH 


ALUMINUM    SALTS.  117 

By  adding  sufficient  ammonium  salt  to  the  solution,  however,  am- 
monium aluminate  and  a  sodium  salt  are  formed,  and  A1(OH)3  is  then 
precipitated. 

Na3A103  +  3NH4Cl-»  (NH4)3A103  +  3NaCl-» 

A1(OH)3  +  MHZ  +  3H20  +  3NaCl 


Cobalt  meta-aluminate  is  probably  the  compound  formed  by  heating 
CoO  with  A1203.  It  is  blue  in  color.  Aluminates  of  calcium  are  among 
the  compounds  formed  when  Portland  cement  "sets." 

In  the  presence  of.tartaric,  citric,  malic  and  other  acids,  sugars, 
starches  and  organic  compounds  containing  the  hydroxyl  radicle, 

A1(OH)3  is  not  precipitated,  owing  to  the  formation  of  compounds  which 

+++ 
produce  stable  complex  ions  and  not  the  simple  Al  ion.    Hence  organic 

matter  in  general  must  be  removed,  usually  by  ignition,  before  aluminum 
can  be  precipitated  as  the  hydroxid.  (See  also  pages  193  and  224.) 

The  carbonate  ion,  C03,  precipitates  A1(OH)3,  this  product  being 
formed  by  complete  hydrolysis  of  the  carbonate  which  is  probably  the 
substance  first  formed. 

The  sulfid  ion,  S,  produces  no  precipitate  in  acid  solution,  for  A12S3, 
formed  only  in  the  '  '  dry  way,  '  '  is  soluble  in  acids  and  is  completely  hy- 
drolyzed  by  water  into  the  hydroxid  and  H2S.  Hence  on  introducing  the 
S  ion  into  solutions  alkaline  with  NH4OH,  aluminum  hydroxid  is  pre- 
cipitated. 

The  sulfate  ion,  S04,  produces  no  precipitate,  for  the  sulfate, 
A12(S04)3,  is  readily  soluble.  This  salt,  like  all  the  common  soluble 
salts  of  aluminum,  deliquesces  greatly  and,  therefore,  can  not  be  well 
preserved.  The  double  sulfate  of  aluminum  and  potassium,  A12(S04)3. 
K2S04.24H20,  does  not  deliquesce,  is  very  soluble  (in  about  15  parts 
H20  at  ordinary  temperatures  and  less  than  one-third  part  H20  at  100°) 
and  is  by  far  the  most  important  salt  of  aluminum  in  the  industrial 
world.  It  is  to  be  regarded  as  the  potassium  salt  of  an  aluminum-sulfuric 
acid  and  its  formula  is  probably  more  correctly  written  KA1(S04)2. 
12H20.  In  concentrated  solutions,  it  doubtless  forms  complex  ions 
(  possibly  Al  (  SO4  )  2  )  to  a  slight  extent,  but  these  are  so  unstable  in  dilute 
solutions  that  only  K,  Al  and  S04  ions  are  in  evidence.  Hence  an 

+-H- 

"alum"  solution  gives  all  the  reactions  of  the  Al  ion. 

The  term,  "alum,"  is  used  to  designate  a  class  of  compounds  of 
similar  composition,  all  of  which  crystallize  in  the  same  form  from 
solution.  In  place  of  aluminum,  trivalent  chromium,  iron  and  manga- 
nese may  be  substituted.  Sodium,  ammonium  and  other  univalent 


118  QUALITATIVE    ANALYSIS. 

metals  may  replace  potassium,  while  sulfur  may  be  similarly  substituted 
by  the  rarer  elements,  tellurium  and  selenium. 

The  chlorid  ion,  01,  produces  no  precipitate,  as  aluminum  chlorid, 
like  other  halids,  is  very  soluble.  The  crystallized  salt,  A1C13.6H2O,  will 
dissolve  in  about  one-fourth  of  its  own  weight  of  water.  As  an  illustra- 
tion of  the  strong  tendency  of  aluminum  salts  to  hydrolyze,  it  may  be 
stated  that  a  solution  of  A1C13  will,  on  evaporation,  leave  A1(OII)3 
entirely.  The  equilibrium  in  this  case  is  prevented  by  the  volatility  of 
HC1  which  passes  off  with  the  water  vapor.  With  salts  of  non-volatile 
acids,  the  introduction  into  the  solution  of  some  reagent,  such  as  BaC03, 
that  will  lower  the  concentration  of  the  free  acid  formed  by  hydrolysis, 
is  sufficient  to  cause  the  reaction  to  run  to  completion. 

The  cyanid  ion,  ON,  precipitates  A1(OH)3  because  of  hydrolysis. 
No  cyanids  of  trivalent  metals  are  known. 

The  phosphate  ion,  P04,  precipitates  white  A1P04,  readily  soluble 
in  dilute  mineral  acids.  Like  chromium  and  ferric  phosphates,  it  is 
insoluble  in  acetic  acid  and,  by  this  means,  may  be  separated  from  the 
metals  forming  Groups  III  and  IV  of  the  qualitative  analysis  scheme. 
Like  aluminum  salts  generally,  it  is  readily  soluble  in  the  strong  alkalies, 
forming  alkaline  aluminates  and  phosphates. 

The  acetate  ion,  (C2H302),  produces  no  precipitate,  as  aluminum  ace- 
tate, A1(C2H302)3,  is  very  soluble.  By  boiling  a  dilute  solution  the  salt  is 
about  two-thirds  hydrolyzed,  and  a  precipitate  is  formed  having  approxi- 
mately the  proportions  indicated  by  the  formula, 

OH 
Al— OH 

(C2H302) 

It  is  known  as  a  basic  acetate  (see  also  pages  194  and  226).  On  cooling 
the  solution,  the  precipitate  dissolves  in  the  acetic  acid  formed  by  the 
hydrolysis.  It  is  thus  a  distinctly  reversible  reaction  with  the  equilibrium 
much  modified  by  changes  of  temperature. 

Group  3,  Second  Type. 

Gallium  70;  Indium  115;  Thallium  204.1. 

The  members  of  this  group  are  trivalent  elements,  in  general  resem- 
bling the  other  type,  although  they  possess  rather  more  metallic  proper- 
ties. As  all  of  them  are  very  rare  substances,  they  will  not  be  con- 
sidered in  detail. 


CARBON    GROUP.  119 

Group  4,  First  Type. 

Carbon  12;          Silicon  28.4;          Titanium  48.1;          Zirconium  90.6; 
Cerium  140;         Thorium  232.5. 

This  group,  as  a  whole,  is  of  intermediate  character,  possessing  neither 
metallic  nor  non-metallic  properties  in  a  marked  degree.  Carbon  and 
silicon  only  are  of  common  occurrence  and,  since  they  are  the  members 
of  lowest  atomic  weights,  they  show  rather  more  non-metallic  properties 
than  the  other  rarer  elements.  Typical  compounds  are  colorless  and 
contain  the  element  in  the  tetravalent  condition.  Compounds  in  which 
the  element  exhibits  divalence  are  well  known,  however.  Beginning 
with  this  group,  the  hydrids  become  of  importance.  They  are  neutral  in 
character  and  do  not  unite  with  water.  The  oxids  of  the  lower  members 
are  the  anhydrids  of  very  weak  acids.  The  oxids  of  the  other  elements 
unite  with  water  to  form  hydroxids,  the  higher  ones  acting  as  bases  and 
those  of  medium  atomic  weights  showing  either  acid  or  basic  properties, 
depending  on  the  environment.  Compounds  of  these  elements  are  strik- 
ingly similar  in  composition  and,  bearing  in  mind  constantly  the  normal 
variation  to  be  expected  within  each  group,  in  properties  also. 

CARBON.     C  — 12. 

The  element  exists  in  at  least  three  allotropic  modifications. 

Diamond  is  the  form  obtained  by  crystallization  at  high  temperatures 
under  great  pressure.  It  is  usually  clear  and  colorless,  is  .the  hardest 
known  substance  and  is  very  resistant  to  all  chemical  action. 

Graphite  is  a  soft,  gray-black  substance,  existing  in  the  form  of 
crystalline  scales,  which  is  almost  as  resistant  to  chemical  action  as 
diamond. 

Amorphous  carbon  is  known  in  the  various  forms:  charcoal,  coke, 
gas-carbon,  bone-black,  and  soot  or  lampblack.  All  of  these  bodies  are 
quite  resistant,  but  not  to  the  same  degree  as  diamond  and  graphite. 
None  of  these  allotropic  forms  melt  under  atmospheric  pressure  but, 
when  heated  to  the  highest  temperature  of  the  electric  arc  or  furnace,  it 
is  just  possible  to  make  them  sublime.  All  modifications  may  be  made 
to  unite  with  oxygen,  i.  e.,  burn,  by  being  strongly  heated.  The  great 
ease  with  which  carbon  is  oxidized  makes  it  one  of  the  best  reducing 
agents  for  high-temperature  reactions.  In  the  form  of  coke  or  charcoal, 
it  is  much  used  to  obtain  metals  from  their  ores.  Carbon  monoxid  is  a 
product  of  incomplete  combustion  but  the  dioxid  is  the  final  product., 

Carbon  monoxid,  CO,  is  a  colorless,  poisonous  gas,  but  little  soluble 
in  water.  It  unites  with  oxygen  readily  at  higher  temperatures  form- 


120  QUALITATIVE    ANALYSIS. 

ing  the  dioxid,  thus  acting  as  a  good  reducing  agent.     It  burns  with  a 
blue  flame. 

Carbon  dioxid,  C02,  is  a  colorless  gas  with  a  slight  pungent  odor  and 
feebly  acid  taste.  It  is  half  again  as  heavy  as  air  and  fairly  soluble  in 
water,  one  volume  of  water  dissolving  somewhat  more  than  an  equal 
volume  of  gas.  The  solution  has  a  slight  acid  reaction,  hence  carbon 
dioxid  must  be  regarded  as  an  acid  anhydrid.  It  may  be  removed  en- 
tirely from  solution  by  boiling,  however. 

Carbon  disulfid,  CS2,  is  an  analogous  body  formed  by  direct  union 
of  the  elements  at  a  high  temperature.  It  is  a  heavy  liquid  with  a  slight 
yellow  color,  boiling  at  46°,  and  is  practically  insoluble  in  water.  It  is 
a  good  solvent  for  fats  and  oils  and  also  for  bromin,  iodin  and  some  other 
inorganic  substances.  It  is  quite  volatile  at  ordinary  temperatures  and, 
as  its  vapors  form  an  explosive  mixture  with  air,  it  should  never  be  used 
near  a  flame  or  heated  substance. 

The  compounds  of  carbon  show  that  it  acts  as  a  weak  acid-forming 
element,  being  for  the  most  part  tetravalent,  although  in  certain  com- 
pounds, which  are  more  or  less  unstable  and  oxidize  readily,  it  appears 
divalent.  It  unites  with  most  metals  to  form  carbids,  many  of  which 
are  decomposed  by  water  or  dilute  acids  with  the  liberation  of  hydrids. 
The  chief  characteristic  of  carbon  is  the  almost  limitless  number  of  com- 
pounds which  it  forms  with  hydrogen,  oxygen  and  a  few  other  elements. 
These  compounds  constitute  the  field  of  organic  chemistry  which  has 
come  to  mean  simply  the  chemistry  of  the  carbon  compounds.  The  vari- 
ation of  properties  found  among  these  bodies  is  almost  as  great  as  that 
encountered  in  the  inorganic  world.  Organic  compounds  are  generally 
less  stable,  however,  and,  on  heating  out  of  contact  with  air,  they  decom- 
pose quite  generally  into  carbon  and  other  substances.  When  heated  in 
the  air,  they  burn  to  C02,  H20  and  other  oxidation  products. 

Compounds  of  carbon  with  hydrogen  only  are  of  several  distinct 
classes,  of  which  ethane,  C2H6,  ethylene,  C2H4,  acetylene,  C2H2,  benzene, 
C6H6,  and  naphthalene,  C10H8,  are  the  more  important  types.  They  are 
the  hydrids  of  carbon  commonly  known  as  hydrocarbons.  Certain  com- 
binations or  radicles,  such  as  CH3 — ,  C2H5 — ,  etc.,  unite  with  other  ele- 
ments to  give  compounds  analogous  to  those  of  the  inorganic  realm. 

Organic  oxids,  compounds  of  these  radicles  with  oxygen,  are  known 
as  ethers.  Of  this  class,  ordinary  ether,  C2H5  —  0  —  C2H5,  is  the  most 
important. 

Organic  hydroxids  are  known  as  alcohols.     Of  these,  wood  alcohol, 
CH3OH,  ordinary  or  ethyl  alcohol,  C2H5OH,  and  amyl  alcohol,  Cg 
are  the  most  important. 


CARBON    ACIDS.  121 

Organic  acids  contain  the  group — 

O 

// 

—  C 

OH 

No  acid  of  carbon  is  as  strong  as  the  common  mineral  acids.  Some,  how- 
ever, dissociate  to  a  greater  degree  than  weak  mineral  acids,  such  as  boric 
and  silicic  acids. 

Carbonic,  acetic,  oxalic  and  tartaric  acids  are  most  important  and 
have  the  composition  and  structure  indicated  by  the  formulae : 

0  HO 

OH                           O               //  \        // 

/                            //               C— OH  HO— C— C— OH 

C=0             CH3— C— OH           |  | 

\                                               C— OH  HO— C— C— OH 

OH                                               \  I         \ 

0  HO 

Carbonic  acid.  Acetic  acid.  Oxalic  acid.  Tartaric  acid. 

Organic  salts  are  formed  by  the  interaction  of  an  organic  acid  with 
an  organic  hydroxid  or  alcohol.  They  are  known  as  esters  and  are 
characterized  by  sweet,  agreeable  odors.  Many  of  the  perfumes  and 
flavoring  extracts  in  general  use  belong  to  this  class  of  compounds. 

The  acid  compounds  of  carbon  are  of  chief  interest  in  analytical 
chemistry  and  the  most  important  of  these  will  be  considered  in  greater 
detail. 

Carbonic  Acid,  H2C03. 

Of  the  more  common  acids  of  carbon,  carbonic  acid  is  the  weakest  and 
most  unstable.  It  has  never  been  obtained  free,  since  it  immediately 
decomposes  into  water  and  the  anhydrid,  C02.  That  it  does  exist  to  a 
slight  extent  in  aqueous  solution,  however,  is  indicated  by  the  fact  that 
carbon  dioxid,  passed  into  water,  gives  to  it  a  slight  acid  reaction. 
Toward  metallic  hydroxids,  it  acts  as  a  dibasic  acid,  forming  salts  of 
most  of  the  metals,  carbonates  of  arsenic,  antimony,  tin,  chromium,  alu- 
minum and  ferric  iron  alone  being  rare  or  unknown.  Although  more 
stable  than  the  free  acid,  with  the  exception  of  the  salts  of  the  alkali 
metals,  all  carbonates  are  decomposed  on  heating  into  C02  and  an  oxid 
of  the  metal.  Carbonates  of  the  alkaline  earths  are  most  difficult  to 
decompose,  in  this  as  in  other  ways  showing  their  close  relation  to  the 
alkali  metals.  Ammonium  carbonate  is  volatile. 


122  QUALITATIVE    ANALYSIS. 

All  normal  carbonates,  with  the  exception  of  those  of  the  alkali  met- 
als and  ammonium,  are  comparatively  insoluble  in  water,  but  are  readily 
decomposed  by  dilute  acids.  On  treatment  of  a  carbonate  with  any 
stronger  acid,  carbonic  acid  is  set  free  and  breaks  down  to  a  large  extent 
into  the  anhydrid,  CO2,  which  passes  out  of  solution.  On  boiling,  the 
decomposition  is  completed  and  all  carbon  dioxid  may  be  removed. 

If  CO2  gas  is  passed  into  water  containing  insoluble  carbonates  of 
the  elements  of  the  alkaline-earth  group  (Ba,  Sr,  Ca,  Mg)  held  in  sus- 
pension, the  precipitate  dissolves,  probably  forming  acid  carbonates  of 
the  type  CaH2(CO3)2  (see  page  95).  These  acid  salts  are  soluble  but 
are  decomposed,  on  boiling,  into  the  normal  carbonate,  water  and  carbon 
dioxid.  The  only  acid  carbonates  known  in  the  solid  state  are  those  of 

the  alkali  metals. 

++ 

The  barium  ion,  Ba,  when  brought  into  contact  with  the  carbonate 
ion,  C03,  gives  a  floccy  white  precipitate,  BaC03,  which,  on  long  standing 
or  boiling,  becomes  crystalline.  (See  page  99.) 

The  silver  ion,  Ag,  precipitates  yellowish-wrhite  Ag2C03.  Like  most 
silver  salts,  it  is  readily  soluble  in  NH4OH.  Like  all  carbonates,  it  dis- 
solves readily  in  HN03. 

Acetic  Acid,  H(C2HS02). 

Acetic  acid  is  a  well-known  monobasic  acid  of  carbon.  Since  it  is 
by  no  means  as  strong  as  the  ordinary  mineral  acids,  it  may  be  set  free 
from  its  salts  by  them.  It  mixes  with  water  in  all  proportions  and  boils 
at  118°.  Its  salts,  like  the  salts  of  organic  acids  generally,  are  decom- 
posed by  heating  into  carbonates.  Hence,  on  ignition,  the  free  acid  may 
be  volatilized  and  its  salts  destroyed.  Fusion  with  ammonium  nitrate 
oxidizes  it  completely. 

All  normal  acetates  are  soluble  in  water,  and  are  transposed  by 
boiling  with  sodium  carbonate. 

The  common  anion  is  (C2H302). 

The  barium  ion,  Ba,  produces  no  precipitate. 

The  silver  ion,  Ag,  produces  no  precipitate  except  in  concentrated 
solutions  when  white  Ag(C2H302)  falls.  It  is  soluble  in  about  100 
parts  of  H20  at  ordinary  temperature. 

-t~H- 

The  ferric  ion,  Fe,  added  to  a  neutral  solution  of  an  acetate  produces 
a  deep  red  color,  due  to  the  formation  of  ferric  acetate,  Fe(C2H302)3, 
which,  on  abundant  dilution  and  boiling,  hydrolyzes  partially,  forming 
a  precipitate  of  basic  ferric  acetate,  Fe(C2H302)  (OH)2,  indistinguish- 
able in  appearance  from  ferric  hydroxid.  (See  page  226.) 


OXALATES.  123 

Ethyl  alcohol  unites  with  acetic  acid  to  form  ethyl  acetate  (see  page 
23),  a  compound  characterized  by  a  peculiar  sweet,  ethereal  odor.  With 
amyl  alcohol  a  similar  compound  is  formed  which  has  the  familiar 
banana  odor.  The  formation  of  these  compounds  takes  place  most 
satisfactorily  in  the  presence  of  concentrated  sulfuric  acid  which  should 
be  warmed.  Heating  should  be  avoided,  as  hot  concentrated  sulfuric 
acid  acts  as  an  oxidizing  agent  (see  page  187)  and  the  sulfur  dioxid 
formed  will  make  the  ethereal  odor  very  difficult  to  detect. 

Oxalic  Acid,  H2C204. 

Oxalic  acid  is  a  common  dibasic  acid  of  carbon,  readily  formed  by 
the  oxidation  of  organic  matter  with  nitric  acid.  Since  all  organic  acids 
are  of  much  the  same  nature,  it  suggests  acetic  acid  in  some  of  its  reac- 
tions. It  is  a  white  solid,  readily  soluble  in  water,  from  which  it  crys- 
tallizes with  two  molecules  of  water  of  crystallization.  It  is  the  strongest 
of  the  common  acids  of  carbon,  dissociating  more  than  phosphoric  but 
less  than  sulfuric  acids  (see  page  16).  On  heating,  it  does  not  melt  but 
sublimes  readily.  Its  salts  are  generally  not  volatile  and,  when  heated, 
they  break  down  in  carbonates.  Hence,  by  ignition,  the  free  acid  may  be 
removed  and  its  salts  destroyed. 

Concentrated  H2S04,  on  heating,  decomposes  oxalic  acid  into  water 
and  equal  volumes  of  carbon  monoxid  and  dioxid. 

H2C204  ->  H20  +  CO  +  C02 

Oxalic  acid  may  be  readily  oxidized  to  C02  and  water,  by  fusion 
with  ammonium  nitrate  or  by  treating  its  solution  with  a  permanganate. 
Each  molecule  of  oxalic  acid  absorbs  one  atom  of  oxygen  and  the  reac- 
tion is  therefore  equivalent  to  the  oxidation  of  CO  to  C02  (see  page  47). 

All  oxalates  are  fairly  insoluble  with  the  exception  of  the  salts  of 
the  alkali  metals  and  of  magnesium.  They  are  readily  soluble  in  dilute 
mineral  acids,  but  not,  as  a  rule,  in  acetic  acid.  In  the  presence  of  an 
excess  of  oxalic  acid  or  an  alkaline  oxalate,  soluble  complex  salts  tend 
to  form  in  many  instances. 

All  oxalates  are  readily  transposed  by  sodium  carbonate,  with  the 
exception  of  salts  of  cobalt,  nickel,  manganese  and  zinc,  which  require 
long  boiling. 

The  common  ion  is  C704. 

The  calcium  ion,  Ca,  causes  a  fine  white  precipitate  of  CaC2O4.2H20 
which  is  exceedingly  insoluble  in  water  and  dilute  acetic  acid.  Being 
more  insoluble  (see  page  97)  than  any  other  salt  of  calcium  that  will 
fall  in  the  presence  of  acetic  acid  (see  page  261),  calcium  oxalate  will  be 


124  QUALITATIVE    ANALYSIS. 

precipitated  when  a  solution  containing  oxalate  ions  is  treated  with  a 
solution  of  calcium  sulfate. 

The  barium  ion,  Ba,  precipitates  white  BaC204,  soluble  in  dilute 
acids,  even  acetic  acid. 

The  silver  ion,  Ag,  precipitates  white,  curdy  silver  oxalate,  Ag2C204, 
soluble  in  HN03  and  in  NH4OH. 

Tartaric  Acid,  H2(C4H408)« 

Tartaric  acid  is  another  well-known,  dibasic  acid  of  carbon.  In  a 
general  way  it  resembles  the  carbon  acids  already  considered.  It  is  a 
colorless  solid,  melting  at  107°,  readily  soluble  in  water  from  which 
it  crystallizes  anhydrous.  Unlike  acetic  and  oxalic  acids,  free  tartaric 
acid  is  not  volatile,  but  is  decomposed  on  heating  with  the  separation  of 
carbon  and  the  production  of  an  odor  like  that  of  "burnt  sugar. "  This 
reaction,  known  as  "charring,"  is  common  to  non-volatile  organic  matter 
in  general,  and,  as  the  phrase  signifies,  the  " burnt  sugar"  odor  is  not 
peculiar  to  tartaric  acid. 

Hot  concentrated  sulfuric  acid  produces  these  same  phenomena  when 
brought  in  contact  with  many  forms  of  organic  matter,  the  concentrated 
acid  aiding  the  decomposition  by  the  abstraction  of  water.  In  addition, 
sulfur  dJoxid  may  be  liberated.  This  gas  is  a  reduction  product  of  sul- 
furic acid  which  gives  up  its  oxygen  to  the  organic  matter  to  form  carbon 
dioxid  and  water,  while  H2S04  becomes  H2S03,  sulfurous  acid.  The 
latter,  being  unstable,  separates  its  anhydrid,  S02,  which  is  volatile. 
Reduction  of  hot  concentrated  sulfuric  acid  is  characteristic  of  organic 
matter  in  general. 

Tartaric  acid,  like  oxalic  acid,  is  readily  oxidized,  either  by  fusion 
with  ammonium  nitrate  or,  in  the  ' '  wet  way, ' '  by  permanganate.  Carbon 
dioxid  and  water  are  the  ultimate  products  of  the  -oxidation  but,  in  the 
"wet  way,"  the  reaction  does  not  under  all  conditions  run  to  the  end. 

Salts  of  tartaric  acid  are  readily  broken  down  on  heating  into  car- 
bonates of  the  metal,  carbon  dioxid,  water  and  free  carbon. 

Tartrates  of  the  alkalies,  except  acid  potassium  and  acid  ammonium 
tartrates  (see  pages  73  and  78),  are  readily  soluble  in  water.  Most 
other  tartrates  are  quite  insoluble  but  are  dissolved  on  the  addition  of 
dilute  acids. 

All  tartrates  are  transposed  by  boiling  with  sodium  carbonate. 

The  common  ion  is  (C4H4O6).  Several  complex  anions  are  recog- 
nized which  contain  within  them  metallic  elements  such  as  copper,  iron, 
aluminum  and  chromium.  (See  pages  81,  224,  117,  193.) 


TAETKATES.  125 

The  barium  ion,  Ba,  added  to  a  neutral  solution  containing  the  tar- 
trate  ion,  produces  a  white  flocculent  precipitate  of  barium  tartrate, 
Ba(C4H406),  soon  becoming  crystalline.  It  is  readily  soluble  in  an  ex- 
cess of  an  alkaline  tartrate,  producing  soluble  double  tartrates.  It  dis- 
solves readily  in  acetic  and  other  acids,  and  in  strong  alkali  solutions 
forming  the  soluble  double  tartrates  just  mentioned. 

-H- 

The  calcium  ion,  Ca,  acts  similarly  to  the  barium  ion.  Although  cal- 
cium tartrate,  Ca(C4H4O6),  requires  more  than  5,000  parts  of  water  for 
its  solution,  at  ordinary  temperatures  it  readily  forms  supersaturated 
solutions,  from  which  the  insoluble  salt  precipitates  only  on  standing. 
Ammonium  salts,  especially  the  chlorid,  tend  to  hold  up  the  precipitate. 

The  silver  ion,  Ag,  precipitates  silver  tartrate,  Ag2(C4H406),  from  a 
neutral  solution.  It  is  readily  soluble  in  dilute  acetic  acid  and  ammo- 
nium hydroxid.  When  the  precipitated  silver  tartrate  is  dissolved  in  a 
few  drops  of  ammonium  hydroxid  and  warmed,  the  salt  decomposes,  pre- 
cipitating metallic  silver  in  the  form  of  a  brilliant  mirror  on  the  sides 
of  the  test-tube. 

Cyanogen,  (CN)2. 

Under  the  influence  of  the  electric  arc,  carbon  and  nitrogen  can  be 
made  to  unite,  forming  cyanogen,  (CN)2,  so  named  because  of  the  deep 
blue  color  of  some  of  the  compounds  it  forms.  Cyanogen  is  a  colorless, 
poisonous  gas  with  an  odor  or  taste  like  * '  bitter  almonds, ' '  burning  with 
a  characteristic  lavender-pink  flame.  It  is  soluble  in  water.  It  unites 
with  yellow  ammonium  sulfid  according  to  the  equation : 

(NH4)2S.S+  (CN)2-»2(NH4)SCN 

In  this  way  it  may  be  changed  into  a  thiocyanate  and  tested  for,  as 
directed  on  page  292. 

Cyanogen  resembles  in  a  general  way  the  less  active  members  of  the 
halogen  group.  Like  chlorin,  it  reacts  with  solutions  of  the  alkalies 
giving  analogous  compounds. 

C12  +  2KOH  ->  KC1  +  KOC1         (CN)  2  +  2KOH  -»  KCN  +  KOCN 

Like  the  halogens,  also,  it  forms  a  compound  with  hydrogen  which  is 
acid  in  nature. 

Hydrocyanic  Acid,  "Prussic  Acid,"  HON. 

Hydrogen  cyanid  is  a  colorless  liquid  with  an  odor  like  "bitter  al- 
monds." It  boils  at  26°  and  freezes  to  an  ice-like  solid  which  melts  at 


126  QUALITATIVE    ANALYSIS. 

—  14°.  It  is  readily  soluble  in  water.  Although  showing  unmistakable 
acid  reactions,  it  stands  near  the  extreme  limit  of  the  weak  acids,  disso- 

4- 

ciating  so  little  in  aqueous  solution  that  the  H  ions  can  scarcely  be  de- 
tected by  litmus  paper.  It  is  unstable  in  solution,  uniting  with  water 
to  form  formic  acid,  H(COOH)  and  ammonia,  which  tend  to  neutralize 
each  other. 

HCN  +  2HOH-»H(COOH)  +NH8=HCOO(NHJ 

The  cyanids  are  a  well-characterized  series  of  salts  of  which  the  al- 
kali, alkaline  earth  and  the  higher  form  of  mercury  are  soluble  in  water. 
Some  are  soluble  with  difficulty.  Others  are  so  highly  insoluble  that 
they  are  not  dissolved  by  dilute  mineral  acids,  contrary  to  the  general 
rule  for  salts  of  weak  acids.  Cyanids  of  metals  with  a  valence  greater 
than  two  are  unknown  (except  in  complex  salts). 

All  cyanids  are  decomposed  by  hot  concentrated  sulfuric  acid  with 
the  formation  of  sulfates.  Some  HCN  is  usually  liberated  but  much 
is  decomposed  into  CO  and  NH3  .  according  to  the  reaction: 

Fe(CN)2  +  2H2S04  +  2H20  ->  FeS04  +  (NHJ  2S04  +  2CO 

All  cyanids  are  transposed  by  boiling  with  sodium  carbonate. 

Soluble  cyanids  hydrolyze  to  a  great  extent  in  solution  because  HCN 
is  such  a  weak  acid.  On  standing  in  the  atmosphere,  carbon  dioxid  is 
absorbed  and  carbonates  of  the  metals  result.  '  Furthermore,  decomposi- 
tion of  the  hydrocyanic  acid  formed  by  the  hydrolysis  takes  place,  as 
already  noted. 

On  heating  out  of  contact  with  air,  cyanids  of  the  alkali  metals  fuse 
unchanged,  while  insoluble  cyanids  are  converted  either  into  (1)  cyano- 
gen and  the  metal,  or  (2)  into  a  metallic  carbid  and  nitrogen,  or  (3) 
entirely  decomposed  into  their  elements. 

Cyanids  are  readily  oxidized  and  are  frequently  used  as  reducing 
agents  in  blowpipe  reactions.  Each  molecule  of  cyanid  absorbs  one 
atom  of  oxygen,  as  shown  in  the  equation  : 


In  alkaline  solution,  yellow  ammonium  sulfid  or  other  reagent  containing 
free  sulfur  oxidizes  a  cyanid  to  a  thiocyanate.  Since  the  ferric  ion  gives 
a  deep  red  color  with  the  thfocyanate  ion  due  to  the  formation  of  ferric 
thiocyanate,  Fe(SCN)3,  this  procedure  offers  a  good  test  for  cyanids. 

Many  insoluble  cyanids  are  soluble  in  solutions  of  alkaline  cyanids, 
forming  complex  cyanids.  Since  many  of  these  solutions  respond  but 
slightly  or  not  at  all  to  tests  for  the  heavy  metal  present  in  them,  indi- 
cating a  great  reduction  in  the  concentration  of  the  heavy  metal  ions, 


CYANIDS.  127 

if  not  their  entire  absence,  it  is  incorrect  to  consider  these  solutions  as 
mixtures  of  the  two  cyanids.  New  compounds  are  formed — complex 
cyanids — dissociating  into  complex  ions  containing  cyanogen  united  with 
heavy  metals  in  various  proportions. '  Four  different  types  are  well  rec- 
ognized, the  metal  and  its  valence  being  indicated  by  M-,  M=,  and 
M-  :  K-  (M-C2N2),  K2=(M=C4N4),  Ks=  (M  =  C6N6),  K4  =  (M  =  C6N6). 
These  bodies  dissociate  as  salts  of  mono-,  di-,  tri-  and  tetrabasic  acids 
as  indicated,  the  heavy  metals  being  contained  in  the  complex  anion 
and  not  existing  in  the  form  of  simple  kathions.  These  complex  ions 
possess  different  degrees  of  stability  and  tend  to  decompose,  more  or 
less,  into  simple  ions,  equilibrium  between  the  simple  and  complex  ions 
being  reached  at  different  stages  in  the  decomposition  process. 

The  acids  corresponding  to  the  first  two  of  the  above  types  are  so 
unstable  that  they  decompose  immediately  into  cyanids  of  hydrogen  and 
the  metal.  Consequently,  when  salts  of  these  types  are  acidified,  hydro- 
cyanic acid  is  liberated  and  the  insoluble  cyanid  precipitated.  To  this 
class  belong  K(Ag-C2N2),  K(Au-C2N2),  K(Cu-C2N2),  K2(Hg  = 
C4N4),  K2(Cd  =  C4N4),  K2(Ni  =  C4NJ  and  K2(Zn  =  C4N4).  The  differ- 
ent degrees  of  instability  and  of  decomposition  into  simple  metallic  ions 
are  illustrated  by  the  different  action  of  hydrogen  sulfid  on  these  salts, 
the  precipitation  of  the  sulfids  taking  place  in  proportion  to  the  concen- 
tration of  metallic  ions  present.  From  solutions  of  the  silver,  mercury 
and  cadmium  salts,  the  corresponding  sulfids  are  precipitated  practically 
completely;  zinc  salts  give  only  a  partial  action  while  nickel  salts  give 
no  precipitate  at  all. 

Compounds  of  the  third  and  fourth  types  are  generally  more  stable. 
Although  K4  =  (Mn  =  C6N6)  and  K4  =  (Co  =  C6N6)  are  decomposed,  by 
acidifying,  yet  H3  -(Co  =  C6N6) ,  H3  =  (Fe  - C6N6)  and  H4=  (Fe  =  C6N6) 
are  sufficiently  stable  to  be  isolated  from  solution.  The  complex  iron 
cyanids  are  most  important  and  will  be  considered  in  detail.  (See 
pages  130  and  131.) 

The  simple  cyanid  ion  is  ON,  for  the  reactions  of  which  a  solution 
of  an  alkaline  cyanid  serves  most  satisfactorily. 

The  barium  ion,  Ba,  causes  no  precipitate. 

The  silver  ion,  Ag,  precipitates  yellowish-white  AgCN,  soluble  in 
ammonium  hydroxid  and  alkaline  cyanids,  but  insoluble  in  nitric  acid. 

The  ferrous  ion,  Fe,  when  introduced  into  a  neutral  or  alkaline  solu- 
tion containing  the  CN  ion,  reacts  slowly  in  the  cold,  but  rapidly  on  boil- 
ing, to  form  the  complex  ferrocyanid  ion,  (Fe  =  C6N6),  according  to 
equation : 

FeS04->K4(Fe  =  C6N6)  +  K2SO4 


128  QUALITATIVE    ANALYSIS. 

or  written  ionically 


This  affords  a  very  delicate  means  of  detecting  the  ON  ion  since  the 
ferrocyanid  ion,  when  treated  with  the  ferric  ion,  precipitates  deep-blue 
ferric  ferrocyanid,  Fe4(Fe  =  C6N6)3.  (See  page  294.) 

Cyanic  Acid,  HOCN. 

Subjected  to  the  action  of  oxidizing  agents,  salts  of  hydrocyanic  acid 
take  up  oxygen,  forming  cyanates.  Cyanic  acid,  HOCN,  is  a  liquid  with 
a  very  strong  odor.  It  is  stronger  than  hydrocyanic  acid  but  not  so 
stable,  as  it  cannot  exist  in  contact  with  water,  but  is  broken  up  into 
carbon  dioxid  and  ammonia. 

HOCN  +  H20  ->  NH3  +  C02 

Therefore,  if  dilute  solutions  of  cyanates  are  treated  with  acids,  car- 
bon dioxid  is  the  gas  which  is  for  the  most  part  liberated,  since  ammonia 
unites  with  the  excess  of  acid  and  remains  in  solution.  From  the  fact 
that  the  cyanates  evolve  C02  (possibly  some  undecomposed  HOCN  also) 
when  treated  with  acids,  they  may  be  readily  distinguished  from  the 
cyanids,  which  evolve  HCN.* 

The  salts  of  cyanic  acid,  in  sharp  contrast  to  the  instability  of  the 
free  acid,  are  very  stable  even  on  heating.  The  alkali  and  alkaline  earth 
salts  are  readily  soluble;  the  silver,  copper  and  lead  and  lower  form  of 
mercury  salts  are  highly  insoluble',  other  salts  being  soluble  with  diffi- 
culty. All  cyanates  are  soluble  in  dilute  mineral  acids  and  are  trans- 
posed by  boiling  with  sodium  carbonate. 

When  a  not  too  dilute  solution  of  a  cyanate  is  treated  with  cobalt 
acetate,  a  blue  color  appears,  due  to  the  formation  of  potassium  cobalto- 
cyanate,  K2(Co04C4N4).  In  dilute  solutions,  this  complex  cyanate  sepa- 
rates into  the  simple  cyanates,  and  the  blue  color  vanishes,  to  reappear 
on  the  addition  of  more  cyanate  or  of  alcohol  which  prevents  the  decom- 
position. This  reaction  is  given  by  thiocyanates  as  well,  but  not  by 
cyanids,  nor  by  cyanates  in  the  presence  of  cyanids. 

The  barium  ion,  Ba,  gives  no  precipitate. 

The  silver  ion,  Ag,  precipitates  white,  curdy  silver  cyanate,  AgOCN, 
which,  like  silver  cyanid,  is  soluble  in  ammonium  hydroxid.  Unlike  sil- 
ver cyanid,  it  is  also  soluble  in  dilute  nitric  acid. 

*  Commercial  potassium  cyanid  usually  contains  potassium  cyanate. 


THIOCYANATES.  129 

Thiocyanic  or  Sulfocyanic  Acid,  HSCN. 

Thiocyanic  acid  is  readily  formed  by  the  action  of  sulfur  on  the 
cyanids  of  the  alkali  metals  (see  page  126).  It  is  a  volatile  liquid, 
B.  P.  85°,  in  general  closely  analogous  to  cyanic  acid,  but  is  much  less 
readily  decomposed  by  water.  Its  salts  are  more  stable.  The  thiocyan- 
ates  of  the  alkalies  fuse  unchanged  out  of  contact  with  air,  showing 
various  colors,  yellow,  green,  blue,  as  the  temperature  rises.  On  cooling 
they  become  white  again.  Thiocyanates  of  the  heavy  metals  decompose 
on  fusion  into  sulfates.  Mercurous  thiocyanate  swells  up  tremendously 
on  heating.  Advantage  is  taken  of  this  property  in  making  the  "Pha- 
roah's  serpents"  frequently  found  among  pyrotechnics. 

Both  the  acid  and  its  salts  are  readily  oxidized.  Heated  in  the  air, 
thiocyanates  are  converted  in  cyanates  and  sulfates. 

KSCN  +  202  +  H20  ->  KHS04  +  HOCN 

The  same  reaction  takes  place  in  alkaline  solution  but,  in  the  pres- 
ence of  acid,  sulfuric  and  hydrocyanic  acids  are  the  usual  products. 

In  HSCN  three  acid-forming  elements  are  united  together.  This 
combination  results  in  the  production  of  one  of  the  strongest  known 
acids.  In  solution,  it  is  dissociated  quite  as  completely  as  hydrochloric 
or  nitric  acids. 

Thiocyanates  are  generally  soluble,  with  the  exception  of  the  salts 
of  silver,  lead,  mercury  and  copper,  all  of  which  are  transposed  on  boil- 
ing with  sodium  carbonate. 

The  common  ion  is  SON,  although  complex  ions  are  known. 

The  barium  ion,  Ba,  gives  no  precipitate. 

The  silver  ion,  Ag,  precipitates  silver  thiocyanate,  AgSCN,  a  white, 
curdy  precipitate,  soluble  in  ammonium  hydroxid  but  not  in  nitric  acid. 
It  is  readily  soluble  in  an  excess  of  an  alkaline  thiocyanate,  forming  a 
complex  salt. 

The  ferric  ion,  Fe,  produces  a  deep  red  color  in  dilute  hydrochloric 
acid  solutions,  due  to  the  formation  of  ferric  thiocyanate,  Fe(SCN)3. 
The  color  becomes  more  intense  in  the  presence  of  an  excess  of  either  salt. 
On  addition  of  mercuric  chlorid,  the  color  vanishes,  due  to  the  formation 
of  a  colorless  complex  mercury  salt  of  thiocyanic  and  hydrochloric  acids. 

2Fe(SCN)8  +  6HgCl2->2FeCl3  +  3(Hg(SCN)2.HgCl2) 


10 


130  QUALITATIVE    ANALYSIS. 

Ferrocyanic  Acid,  H4(Fe  -  C6N6). 

Ferrocyanic  acid  is  a  white  crystalline  solid  acting  as  a  fairly  strong 
tetrabasic  acid,  readily  soluble  in  water,  oxidizing  and  decomposing 
slowly  in  solution  on  exposure  to  air.  All  of  its  salts  are  more  stable 
than  the  free  acid.  The  alkali  and  the  alkaline  earth  metals  form  readily 
soluble  salts,  the  solutions  of  which  react  neutral  to  litmus.  Salts  of 
other  metals  are  generally  highly  insoluble  and  do  not  dissolve  in  cold 
dilute  mineral  acids.  At  the  boiling  temperature,  dilute  sulfuric  acid 
of  fair  concentration  slowly  decomposes  them  with  the  liberation  of  hy- 
drocyanic acid.  Hot  concentrated  sulfuric  acid  breaks  down  all  ferro- 
cyanids  and  sets  free  carbon  monoxid,  according  to  the  equation  : 


FeS04  +  2K2S04  +  3(NH4)2S04  +  600 

Some  ferrocyanids  are  not  readily  transposed  by  boiling  with  sodium 
carbonate  and  require  the  stronger  action  of  the  caustic  alkalies. 

Fe2(Fe  =  C6N6)  +  4KOH-*K4(Fe  =  C6N6)  +  2Fe(OH)2 

On  ignition  all  are  decomposed  into  the  simple  cyanids  which  then  act 
individually.  (See  page  126.) 

Ferrocyanids  are  compounds  of  ferrous  iron,  the  ion,  (Fe  =  C6N6), 
containing  an  atom  of  the  metal  in  the  divalent  condition.  Ferrous  iron 
tends  to  pass  into  the  ferric  condition,  in  complex  ions  only  a  little  less 
readily  than  when  existing  as  simple  ions.  Hence  ferrocyanids  are  mild 
reducing  agents  which,  in  acid  solutions,  tend  to  absorb  oxygen  and 
become  ferricyanids  or  compounds  of  ferric  iron. 

2K4(Fe  =  C6N6)  +  0  +  H2S04-»2K3(Fe  E  C6N6)  +K2S04  +  H20 
or  writing  an  ionic  equation, 

2(Fe  =  C6N6)  +  0  +  H->  2(Fe  -  C6N6)  +  2K  +  H20 

The  ferrocyanid  ion,  (Fe  =  C6N6),  is  one  of  the  most  stable  of  all 
complex  ions.  Even  -at  a  great  dilution  there  is  no  appreciable  decom- 
position into  the  simple  ions.  The  presence  of  H  ions,  which  is  sufficient 
to  decompose  most  other  complex  ions,  is  without  noticeable  effect  at 
ordinary  temperatures  and  as  indicated,  it  requires  the  boiling  tempera- 
ture to  cause  decomposition  which  takes  place  according  to  the  following 
reaction  : 

(Fe=C0Ne)->6CN  +  Fe 


FEKEOCYANIDS.  131 

In  acid  solutions  the  ferrocyanid  ion  passes  over  into  the  ferricyanid 
ion  under  the  influence  of  oxidizing  agents,  as  previously  noted. 

++ 

The  barium  ion,  Ba,  with  the  ferrocyanid  ion  gives  no  precipitate. 

The  silver  ion,  Ag,  precipitates  white  Ag4(Fe  =  C6N6),  insoluble  in 
dilute  nitric  acid  and  ammonium  hydroxid. 

When  the  ferrous  ion,  Fe,  is  added  to  a  solution  containing  the  ferro- 
cyanid ion,  (Fe  =  C6N6),  out  of  contact  with  the  air,  a  white  precipitate 
of  potassium  ferrous  ferrocyanid,  K2Fe(Fe  =  C6N6),  or  ferrous  ferro- 
cyanid, Fe2(Fe  =  C6Ne),  forms.  The  ferrous  ion  is  so  easily  oxidized 
that,  under  ordinary  conditions,  these  precipitates  are  of  a  light  blue 
color,  due  to  the  formation  of  a  small  amount  of  ferric  salt. 

The  ferric  ion,  Fe,  precipitates  ferric  ferrocyanid,  "Prussian  blue," 
Fe4(Fe  =  C6N6)3.  (See  page  226.) 

Ferricyanic  Acid,  H3(Fe  =  C6N6). 

Ferricyanic  acid  is  a  tribasic  acid  similar  in  character  to  ferrocyanic 
acid,  but  less  stable.  Its  alkali,  alkaline  earth'  and  ferric  salts  are 
soluble  in  water;  others  are  insoluble  in  water  and  in  cold  dilute  acids. 
Heated  alone  or  with  sulfuric  acid  the  decomposition  is  the  same  as  that 
of  the  ferrocyanids.  Ferricyanids  are  transposed  by  boiling  with  sodium 
carbonate  ;  in  certain  cases  better  with  sodium  hydroxid. 

Ferricyanids  are  compounds  of  ferric  iron  which,  in  the  presence  of 
reducing  agents,  tends  to  pass  into  the  ferrous  condition  and  form 
ferrocyanids.  This  reaction  takes  place  better  in  alkaline  than  in  acid 
solution. 


2K3(Fe  =  C6N6)+2KOH  +  2II-      ->     2K4(Fe  =  C6N6)  +  2H2O 
or  writing  an  ionic  equation, 

2(Fe-C6Ne)+2  OH  +  2H-      -^     2(Fe  =  C6N6)  +  2H2O 


The  ferricyanid  ion,  (Fe  *&  C0N6),  is  far  more  stable  than  complex 
ions  generally,  but  it  decomposes  more  readily  than  the  ferrocyanid  ion, 
(Fe  =  C6N6).  (See  page  13.) 

The  barium  ion,  Ba,  gives  no  precipitate. 

The  silver  ion,  Ag,  precipitates  orange  silver  ferricyanid, 
Ag3(Fe  =  C6N6),  insoluble  in  nitric  acid,  but  soluble  in  ammonium 
hydroxid. 

The  ferrous  ion,  Fc,  precipitates  dark  blue  ferrous  ferricyanid,  Turn- 


132  QUALITATIVE    ANALYSIS. 

bull's  blue,  Fe3(Fe  =  C6N6)2.    It  does  not  dissolve  in  acids,  but  is  soluble 
in  alkalies,  being  transposed  according  to  the  equation  : 

Fe3(Fe  -,  C6N6)2  +  6NaOH^3Fe(OH)2  +  2Na3(Fe  ^  C6N6) 

The  products  of  this  metathesis  interact,  however,  the  ferricyanid 
breaking  down  into  ferrocyanid  and  oxidizing  ferrous  to  ferric  hydroxid. 

The  ferric  ion,  Fe,  gives  a  brownish-green  or  yellow  color  according 
to  the  dilution  of  the  solution. 

SILICON.     Si  —  28.4. 

Elementary  silicon  is  known  in  two  modifications  which  suggest  the 
similar  forms  of  carbon.  Amorphous  silicon  is  a  soft  brown  powder 
which  melts  at  a  high  temperature  and  then  solidifies  to  a  hard  black 
mass  of  brilliant  crystals  which  are  much  more  resistant  than  the  amor- 
phous form.  Neither  modification  is  active  at  ordinary  temperatures. 
but  each  generally  requires  high  heating  to  make  it  enter  into  combina- 
tion with  other  elements.  Heated  in  the  air,  silicon  burns  to  the  dioxid, 
Si02.  Acids  are  without  effect,  with  the  exception  of  hydrofluoric  acid 
in  which  silicon  dissolves  readily,  thus  showing  some  slight  indication 
of  metallic  properties. 


Silicon  fluorid,  as  well  as  other  halogen  compounds,  is  decomposed 
by  water,  however,  thus  indicating  that  silicon  is  acidic  rather  than 
metallic  in  nature.  This  fact  is  further  evidenced  by  the  formation  of 
silicon  acids  and  of  silicids  of  the  metals.  These  bodies,  like  the  hydrids, 
of  which  several  are  known,  are  in  general  analogous  to  the  correspond- 
ing compounds  of  carbon.  In  all  its  compounds,  silicon  acts  as  a  tetra- 
valent  element. 

Silicon  fluorid,  SiF4,  is  a  heavy,  colorless  gas  with  a  very  pungent, 
stifling  odor.  It  unites  with  hydrogen  fluorid  to  form  hydrosilicofluorie 
acid  (see  page  136)  according  to  the  equation: 

SiF4  +  2HF  ->  H2SiF6 

Silicon  fluorid  is  decomposed  by  water  into  hydrofluoric  and  silicic 
acids,  as  indicated  in  the  following  equation,  but  some  of  the  unchanged 
SiF4  reacts  with  HF  produced  during  the  reaction  to  form  some  H2SiF6. 


SILICON    COMPOUNDS.  IBS. 

H— O— H  II— O\ 

//F        H— 0— H  H— 0\\ 

Si         +  -»  4HF  +  Si 

\\F        H— O— H  H— O// 

\F        H— O— H  H— O/ 

Silicon  dioxid,  silica,  Si02,  exists  in  at  least  three  different  modifica- 
tions, one  amorphous  and  two  crystalline.  The  naturally-occurring 
crystalline  form,  quartz,  is  the  best  known,  "rock  crystal"  being  the 
pure  substance.  Small  amounts  of  different  impurities  give  rise  to  many 
different  varieties  which  are  known  as  rose,  yellow,  smoky  and  black 
quartz,  topaz,  onyx  and  amethyst.  White  sand  consists  almost  entirely 
of  finely  broken  quartz.  All  these  forms  are  exceedingly  hard,  insoluble 
and  resistant  to  chemical  agents  (except  hydrofluoric  acid)  and  fuse  only 
at  the  temperature  of  the  oxyhydrogen  blow  pipe,  i.  e.,  about  1800°- 
2000°. 

Tridymite  is  the  second  modification,  possessing  a  different  crystal- 
line form  from  quartz,  but  otherwise  resembling  it  closely.  It  occurs  as 
a  rock  constituent. 

Amorphous  silica  occurs  as  chalcedony,  agate  and  flint.  These  nat- 
urally-occurring forms  do  not  differ  much  from  quartz,  but  as  obtained 
artificially,  silica  is  more  soluble  and  readily  attacked  by  chemical  re- 
agents than  the  crystalline  form.  With  hydrofluoric  acid,  all  compounds 
of  silicon  react  more  readily  than  with  any  other  common  reagent.  All 
forms  of  silica  dissolve  readily,  forming  silicon  fluorid  and  water. 

Silicon  dioxid  is  the  anhydrid  of  silicic  acid,  from  which  it  is  readily 
formed  by  heating.  More  than  one-quarter  of  the  known  crust  of  the 
earth  is  silicon,  existing  combined  as  the  anhydrid  or  as  salts  of  silicic 
acid. 

Silicic  Acid. 

Ortho-silicic  acid,  H4Si04,  is  exceedingly  unstable  under  ordinary 
conditions.  It  may  lose  water  in  many  different  proportions.  From 
one  molecule  of  acid,  one  molecule  of  water  separates,  forming  meta- 
silicic  acid,  H2Si03.  From  two  molecules  of  H4Si04,  by  the  loss  of  one, 
two  and  three  molecules  respectively  of  water,  H6Si207,  H4Si2O6,  and 
H2Si205,  are  derived,  while  from  three  molecules  of  H4Si04,  by  separa- 
tion of  water,  other  partially-dehydrated  acids  are  formed  as  shown  in 
the  accompanying  equations: 


134  QUALITATIVE    ANALYSIS. 

(1)  2H4Si04  —   H20-»H6Si2OT        (5)   3H4SiO4  —  2H20 -»  H8Si3010 

(2)  "        —  2H,0->H4Si2Oe        (6)         ll        —  3H20  -» H6Si309 

(3)  "         _3H,0-^H2Si205        (7)         ll         -4H2O-»H4Si808 

(4)  3H4SiO4-    HjjO-^H^SisOu     (8)  -5H2O^H2Si307 

These  compounds  are  known  as  polysilicic  acids,  numbers  (2)  and 
(6)  obviously  having  the  same  proportionate  composition  as  meta-silicic 
acid.  Whether  all  of  these  acids  have  an  existence  even  in  solution  is 
very  questionable,  since  no  methods  of  differentiating  them  are  known. 
The  water  content  of  silicic  acid  seems  to  depend  only  on  the  temperature 
to  which  it  is  heated,  nothing  definite  being  obtained  except  the  com- 
pletely dehydrated  SiO2. 

The  solubility  of  the  acid  seems  to  correspond  in  a  general  way  with 
the  state  of  hydration,  the  more  highly-hydrated  forms  being  the  most 
soluble.  In  acids  it  is  somewhat  more  soluble  than  in  pure  water,  while 
solutions  of  the  fixed  alkalies,  both  carbonates  and  hydroxids,  have  a 
very  decided  solvent  action  even  on  the  very  resistant  quartz,  when  it  is 
finely  powdered.  Most  of  the  solutions  possess  the  properties  of  pseudo- 
solutions  quite  as  much  as  of  true  solutions. 

Although  the  anhydrid,  SiO2,  is  not  formed  completely  below  a  red 
heat,  a  temperature  of  110-130°  produces  sufficient  dehydration  to  give 
a  product  insoluble  in  water  and  dilute  acids.  This  drying-out  process, 
known  as  desiccation,  is  frequently  used  as  a  means  for  the  separation  of 
silicic  acid  from  solution. 

Silicic  acid  dissociates  so  little  that  its  aqueous  solutions  give  no  evi- 

+ 

dence  of  the  presence  of  II  ions  when  tested  with  red  litmus.  Silicate 
ions  are,  therefore,  practically  unknown.  Soluble  silicates  undoubtedly 
ionize,  but  hydrolysis  takes  place  to  such  an  extent  as  to  remove  the  sili- 
cate ions  almost  entirely  from  solution  as  free  silicic  acid.  Insoluble  sili- 
cates may  be  precipitated  from  solutions  of  soluble  silicates  on  the  addi- 
tion of  certain  metallic  ions,  it  is  true,  yet  the  concentration  of  silicate 
ions  actually  existing  in  solution  at  any  given  moment  must  be  very  small 
indeed.  None  of  the  reactions  by  which  silicic  acid  is  separated  and 
detected  are  typical  ionic  reactions. 

Salts  of  silicic  acid,  both  simple  silicates  and  polysilicates,  are  well- 
known,  many  being  the  common  rock-forming  minerals.  They  are  among 
the  most  stable  substances  at  high  temperatures.  The  simpler  silicates 
are  generally  of  the  type  derived  from  meta-silicic  acid,  H2SiO3.  With 
the  exception  of  the  silicates  of  the  alkali  metals,  all  are  highly  insoluble, 
many  of  them  being  so  little  soluble  that  they  are  not  affected  by  the  ordi- 
nary strong  acids,  contrary  to  the  general  rule  for  salts  of  weak  acids. 


SILICATES.  135 

Soluble  silicates  of  sodium  and  potassium  are  known  as  "water  glass." 
These  solutions  are  hydrolyzed  to  a  very  great  degree  and  react  strongly 
alkaline. 

Ammonium  silicate  is  unknown,  for  it  cannot  be  formed  in  the  "dry 
way ' '  and  attempts  to  produce  it  in  the  *  *  wet  way ' '  result  only  in  a  com- 
pletely hydrolyzed  product.  To  illustrate :  When  a  solution  of  an  alka- 
line silicate  is  treated  with  the  ammonium  salt  of  a  strong  acid,  such  as 
NH4C1,  there  is  evidence  that  a  metathesis  occurs  to  form  an  alkaline 
chlorid  and  ammonium  silicate.  This  latter  compound  hydrolyzes  and 
the  silicic  acid  is  nearly  all  precipitated  from  a  warm  solution,  especially 
on  standing.  If  ammonium  carbonate  is  substituted  for  ammonium 
chlorid,  the  precipitation  is  not  complete,  since  sodium  carbonate  is  a  sol- 
vent for  silicic  acid.  Ammonium  hydroxid  gives  no  precipitate  under 
these  conditions. 

Many  silicates  are  decomposed  by  acids  and,  on  evaporation  of  the 
solution,  the  silicic  acid  separates  in  the  form  of  a  jelly,  which,  at  120°, 
loses  most  of  its  water,  leaving  an  insoluble  residue  like  powdered  glass. 

Most  silicates,  insoluble  in  acids,  are  not  transposed  by  boiling  with 
sodium  carbonate.  They  may  be  decomposed  by  fusing  with  the  dry 
reagent,  however,  forming  sodium  silicate  and  oxids  of  the  metals,  which 
then  readily  dissolve  in  acids.  Fusion  with  ammonium  carbonate  will  not 
decompose  silicates  in  this  way.  If,  however,  a  finely  powdered  silicate, 
intimately  mixed  with  twice  its  weight  of  ammonium  chlorid  and  four 
times  as  much  calcium  carbonate,  be  heated,  decomposition  occurs  prob- 
ably due  to  the  more  energetic  action  of  ammonium  carbonate  freshly 
formed  by  the  interaction  of  NH4C1  and  CaCO3.  This  method  is  some- 
times necessary  in  decomposing  natural  silicates  which  are  to  be  tested 
for  alkali  metals. 

All  silicates  whatsoever  are  decomposed  by  hydrofluoric  acid,  form- 
ing volatile  silicon  fluorid,  which  may  be  driven  off  by  heat. 

CaSi03  +  6HF  -»  CaF2  +  SiF^  +  3H20 

Since  silicon  fluorid  is  decomposed  by  water  into  hydrofluoric  and 
silicic  acids  (see  page  133),  in  order  to  remove  the  silica  entirely  by 
volatilization,  the  water  set  free  in  the  reaction  must  be  removed.  This 
is  most  satisfactorily  accomplished  by  concentrated  sulfuric  acid.  Since 
glass  is  essentially  a  mixture  of  silicates,  all  reactions  with  hydrofluoric 
acid  must  be  performed  in  platinum  vessels.  (See  page  199.) 

Silica,  being  an  acid  anhydrid,  will  not  unite  with  fused  sodium 
meta-phosphate,  and  therefore  remains  as  an  insoluble,  unfused  residue 


136  QUALITATIVE    ANALYSIS. 

floating  about  in  the  hot  bead.  Most  silicates  are  decomposed  by  fusion 
with  sodium  meta-phosphate,  the  metallic  oxid  uniting  with  the  phos- 
phate and  dissolving  in  the  bead,  leaving  the  " skeleton"  of  Si02.  A 
few  silicates  are  not  thus  decomposed,  and  can  not  be  detected  by  this 
reaction,  since  they  fuse  and  mix  with  the  bead  without  leaving  a 
visible  residue. 

Hydrosilicoflnoric  Acid,  H2SiF6. 

Silicon  fluorid  unites  with  hydrofluoric  acid  to  form  a  complex  body 
according  to  the  equation : 

SiF4  +  2HF-»H2SiF6 

This  compound,  H2SiF6,  is  known  as  hydrosilicofluoric  or  hydro- 
fluosilicic  acid.  The  reaction  by  which  it  is  formed  proceeds  most  satis- 
factorily in  concentrated  solutions  of  hydrofluoric  acid.  Since  silicon 
fluorid  reacts  with  water  to  form  hydrofluoric  acid  (and  silicic  acid), 
hydrosilicofluoric  acid  also  is  formed.  The  following  equation  is  often 
written  to  express  this  reaction,  but  it  is  hardly  probable  that  it  shows 
the  true  way  in  which  hydrosilicofluoric  acid  is  formed : 

3SiF4  +  3H20  ->  2H2SiF6  +  H2Si03 

Hydrosilicofluoric  acid  is  about  as  strong  as  sulfuric  or  dichromic 
acid.  It  is  known  only  in  solution,  for,  on  every  attempt  to  isolate  it,  it 
decomposes  into  silicon  fluorid  and  hydrofluoric  acid. 

All  silicofluorids  are  readily  soluble  in  water  with  the  exceptions  of 
the  salts  of  the  alkali  metals  and  of  barium.  Na2SiF6  is  soluble  in 
about  150  parts  of  H2O,  K2SiF6  in  about  825  parts,  and  BaSiF6  in 
about  3500  parts  at  ordinary  temperatures. 

All  silicofluorids  are  decomposed  in  alkaline  solution  into  salts  of 
silicic  and  hydrofluoric  acids,  according  to  the  equation: 

Na2SiF6  +  GNaOH  ->  6NaF  +  Na2Si03  +  3H20 

Ammonium  hydroxid  precipitates  silicic  acid  because  of  hydrolysis 
of  the  ammonium  silicate.  (See  page  135.) 

All  silicofluorids  are  decomposed  on  heating  into  fluorids  of  silicon 
and  of  the  metal.  Concentrated  sulfuric  acid  forms  sulfates  and  liber- 
ates silicon  fluorid  and  hydrofluoric  acid. 

The  silicifluorid  ion  is  SiF«. 


TIN   GROUP.  137 

-H- 

The  barium  ion,  Ba,  precipitates  white  crystalline  BaSiF6  which  is 
somewhat  soluble  in  HC1.     (See  page  100.) 
The  silver  ion,  Ag,  gives  no  precipitate. 

Group  4,  Second  Type. 
Germanium  72.5;          Tin  119;          Lead  206.9. 

The  elements  of  this  type  closely  resemble  those  of  the  first  type, 
but  are  rather  more  metallic  in  properties,  being  soft,  lustrous,  heavy, 
easily-fusible  metals.  The  hydroxids  ordinarily  act  as  weak  bases,  but, 
under  proper  conditions,  may  show  acid  properties.  The  valence  of  the 
element  is  quite  as  often  two  as  four.  The  metals  do  not  pass  readily 
into  the  ionic  conditions,  hence  they  may  be  obtained  from  their  com- 
pounds without  difficulty. 

Germanium  is  of  rare  occurrence  and  will  not  be  treated  further. 

TIN  (Stannum).     Sn  — 119. 

Tin  is  a  rather  soft,  silver-white  metal,  under  ordinary  conditions 
possessing  a  crystalline  structure.  It  does  not  tarnish  in  the  air  at 
ordinary  temperatures.  When  heated,  it  oxidizes  slowly  until  a  white 
heat  is  reached,  when  it  burns  rapidly  to  the  white  oxid,  Sn02. 

The  metal  dissolves  slowly  in  fairly  dilute  solutions  of  the  strong 
mineral  acids,  forming  in  every  instance  stannous  compounds  in  which 
the  metal  is  bivalent.  In  hot  concentrated  acids  the  action  often  differs 
from  that  in  dilute  solutions.  Hydrochloric  acid  forms  stannous  chlorid, 
SnCl2,  but  sulfuric  acid  produces  the  tetravalent  stannic  sulfate, 
Sn(S04)2,  together  with  S02,  S  and  H20. 

Fairly  concentrated  HN03  doubtless  oxidizes  the  stannous  nitrate, 
first  formed,  over  to  stannic  nitrate.  Tetravalent  tin  salts  are  weaker 
than  the  bivalent  compounds  and  hydrolyze  very  considerably.  In  the 
present  instance  stannic  hydroxid,  Sn(OH)4,  is  formed.  This  loses  a 
molecule  of  water  and  becomes  SnO(OH)2,  several  molecules  of  which 
unite  to  form  a  compound  molecule,  (SnO(OH)2)x,  where  x  possibly 
equals  five.  Stannic  hydroxid  possesses  acidic  as  well  as  basic  properties, 
hence  it  is  often  called  stannic  acid.  The  highly  insoluble  compound 
formed  by  the  action  of  concentrated  HNO3  on  tin  is  known  as  /?  stannic 
acid  to  distinguish  it  from  the  ordinary  soluble  form. 

Tin  is  quite  indifferent  to  the  passage  into  the  ionic  condition,  no  con- 
siderable amount  (if  any)  of  energy  being  absorbed  or  liberated  during 
the  reaction.  For  this  reason,  several  metals  showing  a  greater  tendency 
to  pass  into  the  ionic  condition,  especially  zinc,  magnesium  and  aluminum, 


138  QUALITATIVE    ANALYSIS. 

will  precipitate  metallic  tin  from  solutions  of  its  salts.  On  the  other 
hand,  silver,  mercury,  bismuth,  copper,  platinum  and  gold  are  thrown 
out  of  solutions  of  their  salts  when  metallic  tin  is  introduced. 

Stannous  compounds  show  a  very  strong  tendency  to  become  tetra- 
valent  and  will  react  with  most  oxidizing  agents  and  pass  into  the 
stannic  condition.  For  this  reason  stannous  chlorid,  SnCl2,  is  one  of 
the  most  satisfactory  reducing  agents  for  general  laboratory  use.  By  it, 
mercuric  salts  are  changed  first  into  the  mercurous  condition  and  finally 
into  the  metal.  Compounds  of  arsenic  in  both  stages  of  oxidation  are 
reduced  to  the  element  (Bettendorf  Test  for  arsenic).  Ferric  salts  are 
reduced  to  the  ferrous  condition  and  the  highly  oxidized  acidic  compounds 
of  chromium  and  manganese  are  reduced  to  the  basic  condition. 

Stannic  compounds,  on  the  other  hand,  are  reduced  only  with  great 
difficulty,  nascent  hydrogen  being  the  only  agent  capable  of  breaking 
down  stannic  salts  in  solution  into  the  lower  condition. 

Heated  before  the  blowpipe  with  Na2C03  on  charcoal,  all  tin  com- 
pounds are  reduced  to  the  metal.  The  reaction  may  be  accomplished  most 
readily  by  adding  finely  powdered  charcoal  to  the  mixture  of  the  tin 
compound  and  Na2C03.  Even  under  these  conditions  long  and  steady 
blow-piping  is  sometimes  necessary  to  obtain  the  bright,  white  metallic 
globule  which,  as  soon  as  the  flame  is  removed,  becomes  covered  with  a 
coating  of  white  oxid.  The  globule  is  malleable  (distinction  from  Sb 
and  Bi).  It  does  not  dissolve  in  HN03,  but  is  changed  to  a  white  powder 
(like  Sb  but  different  from  Pb  and  Ag).  The  globule  is  readily  soluble 
in  HC1.  A  white  coating  of  the  oxid  is  formed  on  the  charcoal. 

The  compounds  of  tin  in  both  stages  of  oxidation  are  generally 
colorless.  Salts  of  the  common  strong  mineral  acids  are  usually  soluble 
but  since  tin  is  but  a  weak  base-forming  element,  they  tend  to  hydrolyze. 
Hence  a  little  acid  must  be  added  to  the  solutions  of  tin  salts  to  prevent 
the  formation  of  insoluble  basic  salts. 

Of  the  tin  ions,  the  stannous  ion,  Sn,  is  well  known.  Although  the 
stannic  ion,  Sn,  is  recognized,  its  compounds  dissociate  but  little  and  in 
solution  the  tin  exists  for  the  most  part  as  pseudo-solutions  of  stannic 
hydroxid.  Tin  enters  into  several  anions  of  which  the  stannite  ion,  Sn02, 
the  stannate  ion,  Sn03,  and  the  tlyostannate  ion,  SnS3,  are  the  most 
important. 

Since  the  compounds  of  the  Sn  ion  are  quite  different  from  the  com- 
pounds of  the  Sn  ion  in  their  properties,  they  will  be  considered  sepa- 
rately. Solutions  of  the  chlorids  serve  satisfactorily  for  the  reactions 
of  each  ion. 


STANNOUS    SALTS.  139 

The  Stannous  Ion,  Sn. 

The  hydroxyl  ion,  OH,  added  to  a  solution  containing  the  S+n  ion, 
precipitates  white  Sn(OH)2  which  loses  water  on  heating  and  breaks 
down  into  the  black  SnO.  Stannous  oxid  unites  with  oxygen  so  readily, 
however,  that  it  will  take  fire  and  burn  to  white  stannic  oxid,  Sn02. 

In  the  bivalent  Stannous  hydroxid,  basic  properties  are  more  strongly 
developed  than  acid  properties.  The  latter  are  sufficiently  strong,  how- 
ever, to  cause  the  hydroxid  to  dissolve  in  solutions  of  the  strong  alkalies, 
NaOH  and  KOH,  but  when  the  concentration  of  OH  ions  is  much  less  as  in 
solutions  of  NH4OH  and  in  alkaline  carbonates,  stannous  hydroxid  is 
not  appreciably  soluble.  But  few  salts  of  stannous  acid,  H2Sn02,  are 
definitely  known,  and  aside  from  Na2SnO2  and  K2SnO2  they  are  of  little 
importance.  Solutions  of  these  salts  are  very  strongly  hydrolyzed.  This 
action  increases  so  much  as  the  temperature  is  raised  that  a  brown-black 
precipitate  of  SnO  is  formed.  The  reactions  producing  this  body  are 
simply  hydrolysis  and  dehydration,  as  shown  by  the  equations: 

Na2SnO2  +  2H20  ->  Sn(OH)2  +  2NaOH 


Like  all  bivalent  tin  compounds,  alkaline  stannites  are  very  strong 
reducing  agents.  When  heated,  especially  if  KOH  or  NaOH  be  present 
in  excess,  one  portion  of  the  stannite  is  oxidized  to  stannate  at  the  ex- 
pense of  the  other,  which  is  reduced  to  the  metal  and  precipitated  as  a 
black  powder. 

2Na2Sn02  +  H2O  -»  Na2Sn03  +  Sn  +  2NaOH 

The  reducing  action  of  alkaline  stannites  is  used  in  testing  for  bismuth 
salts.  (See  page  171.) 

The  carbonate  ion,  C03,  has  no  specific  effect,  as  stannous  carbonate 
is  unknown.  Alkaline  carbonate  solutions  act  like  very  weak  hydroxid 
solutions,  precipitating  Sn(OH)2. 

The  sulfid  ion,  S,  precipitates  dark  brown  stannous  sulfid  from  acid 
solutions  only.  The  concentration  o|fj&  ions  must  not  be  too  great,  for  SnS 
is  readily  soluble  in  strong  HC1.  J^aline  polysulfids,  because  of  their 
excess  of  sulfur,  oxidize  SnS  over  to  SnS2,  which  then  dissolves  readily, 
forming  alkaline  thiostannates.  In  monosulfids,  SnS  is  not  appreciably 
soluble  and  apparently  thiostannites  do  not  exist.  Stannous  sulfid  dis- 
solves in  strong  alkalies  with  the  absorption  of  oxygen,  forming  a 
mixture  of  stannites  and  thiostannates. 


140  QUALITATIVE    ANALYSIS. 

The  sulfate  ion,  S04,  produces  no  precipitate,  since  SnS04  is  soluble. 

The  chlorid  ion,  01,  produces  no  precipitate  since  stannous  chlorid 
is  soluble  in  about  one-third  its  weight  of  water.  It  crystallizes  according 
to  the  formula  SnCl2.2H20.  Being  by  far  the  most  important  compound, 

it  is  known  commercially  as  1 1  tin  salt. ' '    In  solution  it  tends  to  precipi- 

OTT 
tate  a  white  insoluble  basic  salt,  Sn<  „,    ,  because  of  hydrolysis;  hence 

a  little  HC1  must  he  added  in  order  to  keep  the  solution  clear.  In  order 
to  prevent  oxidation  by  the  air  to  SnCl4,  it  is  necessary  to  keep  metallic 
tin  in  the  solution.  This,  reacting  slowly  with  the  acid,  generates  nescent 
hydrogen  and  preserves  the  tin  in  the  lower  form. 

The  iodid  ion,  I,  precipitates  orange-brown  SnI2,  which  separates  in 
gleaming  spangles.  It  is  quite  soluble  in  water,  hence  is  not  precipitated 
from  dilute  solutions.  It  dissolves  in  an  excess  of  the  precipitant,  form- 
ing a  double  iodid. 

The  cyanid  ion,  CN,  forms  probably  stannous  cyanid,  but  this  is  com- 
pletely hydrolyzed  and  stannous  hydroxid  is  precipitated. 

The  phosphate  ion,  P04,  causes  the  precipitation  of  a  white  oxyphos- 
phate,  of  varying  composition,  depending  upon  the  conditions  of  the  ex- 
periment and  more  or  less  closely  approximating  the  proportions  indi- 
cated by  the  formula  Sn3(PO4)2.  It  is  soluble  in  strong  acids,  and  in 
strong  alkalies  forming  alkaline  stannites  and  phosphates. 

The  Stannic  Ion,  Sn. 

The  hydroxyl  ion,  OH,  added  to  solutions  containing  ^e  Sn  ion, 
probably  forms  Sn(OH)4  at  first,  but  this  loses  water  and  wrhite 
SnO(OH)2  is  precipitated.  On  heating,  white  Sn02  is  formed.  Tin 
dioxid  is  more  resistant  to  the  odinary  chemical  reagents  than  most 
substances  of  frequent  occurrence.  Neither  acids  nor  aqua  regia, 
nor  fusion  with  alkaline  hydroxids  or  carbonates  serve  to  convert  it 
entirely  into  a  soluble  product.  By  fusing  with  a  mixture  of  sulfur  and 
alkaline  hydroxid  or  carbonate,  stannic  oxid  is  changed  into  soluble 
thiostannates,  however. 

Stannic  hydroxid  possesses  less  basic  properties  than  stannous  hy- 
droxid but  correspondingly  stronger  acid  properties.  It  dissolves  in  acids 
to  form  stannic  salts  and  in  the  strong  alkalies  to  form  alkaline  stannates, 
which  form  the  ion  Sn03.  Since  it  is  rather  acidic  than  basic,  H2SnO3 
is  frequently  called  stannic  acid.  Its  formula  shows  that  it  is  of  the 
same  type  as  the  more  unstable  carbonic  acid,  H2CO3.  The  fact  that 
stannic  acid  anhydrid,  SnO2,  is  a  very  resistant  solid  while  carbonic  acid 


STANNIC    SALTS.  141 

anhydrid,  C02,  is  volatile,  doubtless  accounts  to  a  great  degreee  for  the 
difference  in  stability  between  the  two  acids  and  their  salts.  The  com- 
pound precipitated  from  solutions  of  stannic  salts  by  the  hydroxyl  ion 
seems  to  have  a  single  molecule,  H2Sn03,  and  is  far  more  soluble  and 
reactive  than  the  polymeric  compound  formed  by  the  action  of  fairly 
concentrated  HN03  on  the  metal.  The  former  is  known  as  a  stannic  acid 
to  distinguish  it  from  the  latter  compound. 

The  carbonate  ion,  C03,  has  no  specific  action.  Solutions  of  alkaline 
carbonates  precipitate  H2Sn03  because  of  their  hydrolysis. 

The  sulfid  ion,  S,  precipitates  yellow  stannic  sulfid,  SnS2,  from  acid 
solutions  only.  The  concentration  of  H  ions  must  not  be  too  great,  how- 
ever, for  SnS2  is  readily  soluble  in  fairly  concentrated  HC1.  In  the 
strong  alkalies,  alkaline  sulfids  and  polysulfids,  it  is  soluble,  forming  thio- 
stannates.  In  NH4OH  and  (NH4)2C03,  it  is  insoluble  even  when  heated 
(distinction  from  As). 

The  sulfate  ion,  S04,  causes  no  precipitate,  as  Sn(S04)2.2H20  is  solu- 
ble in  water  containing  sufficient  acid  to  prevent  hydrolysis. 

The  chlorid  ion,  Cl,  produces  no  precipitate,  as  SnCl4  is  very  soluble. 
The  solutions  hydrolyze  considerably  and,  on  standing,  precipitate  white 
basic  salts.  On  boiling  a  dilute  SnCl4  solution,  complete  hydrolysis  and 
dehydration  of  the  resulting  stannic  hydroxid  takes  place  and  stannic 
oxid  is  precipitated. 

The  other  halids  are  similar  to  the  chlorid. 

The  cyanid  ion,  ON,  precipitates  stannic  hydroxid. 

The  phosphate  ion,  P04,  precipitates  white  stannic  oxyphosphate, 
which  is  soluble  in  the  strong  acids  and  alkalies.  A  formula  for  this 
compound  may  be  written : 

Sn^P04 

0 
Sn  =  P04 

Thiostannates  bear  the  same  relation  to  stannates  that  thioarsen- 
ates  and  thioantimonates  do  to  arsenates  and  antimonates  (see  pages  164 
and  169).  Although  thioarsenites  and  thioantimonites  are  well  known, 
the  corresponding  thiostannites  do  not  seem  to  exist.  The  methods  of 
formation  of  all  salts  of  thioacids  are  in  general  alike.  Thiostannates 
are  produced  by  the  action  of  alkaline  polysulfids  on  SnS,  and  by  alka- 
lies and  alkaline  sulfids  on  SnS2,  as  indicated  by  the  equations : 


142  QUALITATIVE    ANALYSIS. 

SnS+  (NHJ2S.xS-»  (NHJ2SnS3+  (x 
3SnS2  +  6KOH  -*  2K2SnS3  +  K2Sn03  +  3H2O 

SnS2  +  (NHJ2S  ->  (NH4)2SnS3 

Aside  from  the  alkaline  thiostannates,  few  salts  are  known.  The  acid 
is  not  known  with  certainty,  for  it  is  very  unstable  and  breaks  down  into 
H2S  and  SnS  analogously  to  the  corresponding  thioarsenic  and  thioanti- 
monic  acids.  Consequently,  on  acidifying  a  solution  of  a  thiostann.ate, 
yellow  SnS2  is  precipitated. 

LEAD  (Plumbum) .     Pb  —  206.9. 

Lead  is  a  heavy,  blue-gray  metal  with  considerable  lustre,  quickly  tar- 
nishing in  the  air  because  of  the  formation  of  a  thin  film  of  Pb20,  a  lower 
oxid  of  little  relative  importance.  At  a  higher  temperature  this  com- 
pound, as  well  as  the  metal,  is  changed  into  PbO  or  litharge. 

The  metal  is  but  little  acted  upon  by  cold  dilute  HC1  or  H2S04,  owing 
to  the  formation  of  a  layer  of  insoluble  chlorid  or  sulfate.  In  hot  HC1 
or  concentrated  H2S04  it  dissolves,  however,  since  the  protective  coatings 
are  soluble  under  these  conditions.  Nitric  acid  and  acetic  acid  are  ready 
solvents  for  the  metal,  forming  the  only  common  soluble  salts.  Lead 
compounds  are  generally  quite  insoluble  and  are,  for  the  most  part, 
colorless. 

Lead  is  indifferent  toward  the  passage  from  the  metallic  into  the  ionic 
state.  No  considerable  amount  of  energy  is  either  liberated  or  absorbed 
when  the  metal  becomes  the  ion  or  when  the  reverse  process  takes  place. 
Metallic  magnesium,  zinc  and  iron,  because  of  their  strong  tendency  to 
become  ionic,  will  precipitate  metallic  lead  from  solutions  of  its  salts,  but 
metallic  lead  will  precipitate  mercury,  silver,  gold  and  platinum. 

On  heating  before  the  blowpipe  with  Na2CO3  on  charcoal,  all  lead  salts 
are  reduced  to  the  metal,  forming  a  bright,  malleable  (difference  from  Sb 
and  Bi)  globule  with  a  yellow  coating  on  the  charcoal  (difference  from 
Sn,  Ag  and  Bi).  It  dissolves  readily  in  HN03  (difference  from  Sn)  and 
gives  a  precipitate  with  HC1  only  in  cold,  fairly  concentrated  solutions 
(difference  from  Ag). 

From  the  two  typical  oxids,  PbO  and  Pb02,  two  different  series  of 
compounds  are  formed.  In  the  divalent  condition  lead  is  generally  basic 
and  from  the  hydroxid,  Pb(OH)2,  all  the  common  salts  are  derived. 
Tetravalent  lead  is  almost  entirely  acidic  in  its  properties.  The  tetra- 
valent  hydroxid  is  a  weak  acid,  of  which  two  forms  are  known : 


LEAD    COMPOUNDS.  143 

/OH  OH 

//OH  / 

Pb  Ortho-plumbic  acid.  0  =  Pb  Meta-plumbic  acid. 

\\OH  \ 

\OH  OH 

Well-characterized  salts  of  these  acids  are  known,  among  which  the  oxids, 
Pb304  and  Pb203,  are  to  be  included.  Minium,  or  "red  lead,"  is  to  be 
regarded  as  the  lead  salt  of  ortho-plumbic  acid,  while  Pb2O3  is  the  lead 
salt  of  meta-plumbic  acid.  This  is  indicated  by  the  fact  that  when  Pb304 
is  treated  with  dilute  HN03,  two-thirds  of  the  lead  dissolves  while  one- 
third  remains  as  insoluble  Pb02,  the  anhydrid  of  ortho-plumbic  acid. 
(See  also  page  214.) 


/o\ 

P(b  O  =  Pb<n>Pb 

\\0\ 
\0/Pb 

The  tetravalent  salts  of  lead  are  of  little  relative  importance  in  qualita- 
tive analysis. 

The  common  ion  is  Pb.  Lead  is  a  constituent  of  some  anions  of  which 
Pb02  is  the  best  known. 

For  reactions  of  the  Pb  ion,  solutions  of  the  nitrate  or  acetate  are  very 
satisfactory.  The  nitrate,  Pb(N03)2,  dissolves  in  about  two  parts  of 
water  and  separates  from  solution  without  water  of  crystallization. 
"  Sugar  of  lead,"  Pb(C2H,O2)2.3H20,  is  about  as  soluble  but,  on  expos- 
ure to  the  air,  becomes  turbid  owing  to  the  formation  of  a  white  basic  car- 
bonate, due  to  the  action  of  the  carbon  dioxid  present  in  the  atmosphere. 

The  hydroxyl  ion,  OH,  precipitates  from  solutions  containing  the  Pb 
ion  white,  floccy  Pb(OH)2.  After  being  washed  free  from  all  of  the 
reagent  used  to  precipitate  it,  lead  hydroxid  gives  a  slight  basic  reaction 
to  litmus,  thus  showing  it  to  be  somewhat  soluble  (8,000  parts  H20)  and 
to  separate  OH  ions  to  a  slight  degree. 

Lead  hydroxid  is  also  capable  of  separating  H  ions  and  of  acting  as 
a  very  weak  acid.  A  high  concentration  of  hydroxyl  ions,  such  as  is 
given  by  the  strong  alkaline  hydroxids,  is  necessary  to  cause  this  change 
in  character.  When  NaOH  is  added  to  a  solution  containing  Pb02H2 

in  suspension,  the  readily-dissociating  strong  base  represses  the  ionization 

+f 

of  lead  hydroxid  as  a  base  (see  page  31).  The  concentration  of  Pb 
ions  is  thereby  lessened  and  Pb02H2  passes  into  solution  (see  page 


144  QUALITATIVE    ANALYSIS. 

36).  On  the  other  hand,  the  tendency  for  lead  hydroxid  to  separate 
H  and  Pb02  ions  is  stimulated,  owing  to  the  fact  that  the  OH  ions  from 

+ 

the  alkali  unite  with  the  H  ions  from  the  Pb02H2  to  form  undissociated 
water.  The  only  lead-containing  ion  left  in  solution  by  these  reactions 
is  the  plumbite  ion,  Pb02,  together  with  an  abundance  of  ions  of  the 
alkali  metal  to  form,  on  evaporation  of  the  solution,  an  alkaline  plumbite. 
Thus,  Pb02H2  may  be  either  the  base — lead  hydroxid — or  plumbous  acid. 

In  the  presence  of  NH4OH,  lead  hydroxid  is  insoluble,  since  the  for- 
mer is  not  sufficiently  strong  to  make  the  latter  act  as  plumbus  acid. 

With  the  exception  of  a  few  very  highly  insoluble  compounds,  all 
lead  salts  are  appreciably  soluble  in  the  strong  alkaline  hydroxids. 

The  carbonate  ion,  C03,  precipitates  white  PbC03.  As  ordinarily 
brought  down  by  alkaline  carbonates,  however,  the  precipitate  consists 
partly  of  Pb(OH)2.  Because  of  the  hydrolysis  of  the  precipitants,  their 
solutions  contain  the  OH  ion  as  well  as  the  C03  ion.  Since  PbC03  and 
Pb(OH)2  are  about  equally  soluble,  a  mixture  of  these  two  salts  falls. 
Whether  this  precipitate,  the  ' '  white  lead ' '  of  commerce,  is  really  a  mix- 
ture or  a  true  basic  salt  (see  page  21)  is  not  yet  definitely  proven. 

The  sulfid  ion,  S,  precipitates  black  PbS  from  both  alkaline  and  acid 
solutions.  From  fairly  acid  solutions,  especially  in  the  presence  of  HC1, 
H2S  precipitates  a  red  complex  salt  having  the  proportions  represented 
by  the  formula  PbS.PbCl2.  On  further  treatment  with  H2S,  the  normal 
sulfid  is  formed.  Lead  sulfid  is  soluble  in  strong  HC1  and  in  hot,  dilute 
HNO3,  since  the  boiling  reagent  oxidizes  it  to  free  sulfur  (or  H2S04)  and 
Pb(N03)2. 

The  sulf ate  ion,  S04,  causes  the  precipitation  of  granular,  white 
PbS04  which  is  soluble  in  about  23,000  parts  of  water  at  ordinary  tem- 
peratures and  somewhat  more  soluble  in  dilute  solutions  of  the  stronger 
acids.  In  concentrated  H2S04,  it  dissolves  to  a  considerable  extent,  ow- 
ing to  the  formation  of  a  soluble  acid  salt.  On  dilution,  however,  the  acid 
salt  is  decomposed  and  PbS04  again  precipitated.  Commercial  H2S04 
always  contains  some  of  this  salt,  since  the  acid  is  concentrated  in  lead 
pans.  Lead  sulfate  is  soluble  in  NaOH  and  KOH,  forming  alkaline 
plumbites  and  sulfates  both  of  which  are  soluble.  Furthermore,  it  dis- 
solves in  strong  solutions  of  the  ammonium  salts  of  several  organic  acids, 
the  acetate,  tartrate  and  citrate  being  most  commonly  used.  The  reac- 
tion  with  ammonium  acetate  appears  to  be  simply  a  metathesis  forming 
lead  acetate.  This  substance,  normally  dissociating  much  less  than  other 
^  salts,  has  its  ionization  repressed  by  the  greater  dissociation  of  ammonium 


NITEOGEN    GKOUP.  145 

++ 
acetate.     Therefore  but  few  Pb  ions  exist  in  solution,  the  concentration 

being  insufficient  to  precipitate  PbS04  again. 

The  chromate  ion,  Cr04,  precipitates  yellow  PbCr04,  soluble  in  about 
100,000  parts  of  water  at  ordinary  temperatures  and  in  the  strong  acids 
and  alkalies  as  would  be  expected.  (See  opposite  page.) 

The   chlorid   ion,    Cl,    precipitates   curdy,   white   PbCl2,    soluble   in 

about  100  parts  of  water  at  ordinary  temperatures  and  in  about  25  parts 

++ 

of  boiling  pater.    For  this  reason,  the  Pb  ion  can  be  removed  only  par- 
tially from  solution  by  the  addition  of  the  Cl  ion. 

The  bromid  ion,  Br,  precipitates  pale  yellow  PbBr2,  similar  to  PbCl2, 
but  less  soluble. 

The  iodid  ion,  I,  brings  down  yellow  PbI2,  similar  to  the  other  halids 
of  lead,  but  less  soluble,  dissolving  in  about  1300  parts  of  water  at  ordi- 
nary temperatures  and  194  parts  at  212°.  On  cooling,  the  salt  separates 
in  the  form  of  brilliant  spangles.  Lead  iodid  is  appreciably  soluble  in 
an  excess  of  KI  forming  a  double  salt  of  the  formula  PbI2.2KI.  In 
strong  solutions  the  ions  are  2K  and  PbI4,  thus  showing  the  formation 
of  a  new  compound  in  which  the  lead  is  found  in  the  anion.  This 
tendency  toward  the  formation  of  new  ions  of  this  character  is  noticeable 
with  the  bromids  and  not  very  appreciable  among  the  chlorids.  With 
the  iodids,  however,  it  is  well  developed  in  the  salts  of  other  metals  as 
well  as  those  of  lead.  (See  similar  mercury  salts,  page  109.) 

The  cyanid  ion,  ON,  precipitates  white  Pb(CN)2,  insoluble  in  an  ex- 
cess of  alkaline  cyanids,  but  dissolving  in  acids  and  strong  alkalies. 

The  phosphate  ion,  P04,  brings  down  floccy  white  Pb3(P04)2,  readily 
soluble  in  acids  and  in  strong  alkalies. 

Group  5,  First  Type. 

Nitrogen  14 ;         Phosphorus  31 ;         Arsenic  75 ;     .    Antimony  120.2 ; 

Bismuth  208.5. 

The  members  of  this  group  show  characteristic  valences  of  three  or 
five.  The  elements  with  low  atomic  weights,  nitrogen  and  phosphorus, 
are  typical  non-metals,  the  pentavalent  compounds  of  which  are  more 
stable  than  the  trivalent.  The  elements  of  high  atomic  weights  are  typic- 
ally metallic  in  appearance,  being  heavy  and  lustrous.  Pentavalent 
compounds  of  bismuth  and  antimony  are  less  stable  than  the  correspon- 
ding trivalent  compounds,  while  with  arsenic  each  form  possesses  ap- 
proximately equal  stability.  The  hydrids  are  rather  unstable,  those  of 
nitrogen  and  phosphorus  uniting  with  water  to  form  weak  bases,  while 
those  of  arsenic  and  antimony  are  neutral  in  character. 

11 


146  QUALITATIVE    ANALYSIS. 

Oxygen  unites  with  these  elements  readily  in  various  proportions, 
trioxids  and  pentoxids  being  most  typical.  These  unite  with  water, 
generally  directly,  to  form  hydroxids.  The  gradation  of  properties 
noticed  in  all  groups  is  especially  prominent  in  this  group.  Trivalent 
hydroxids  of  the  elements  with  low  atomic  weights  are  acids,  while  the 
corresponding  compounds  of  elements  with  high  atomic  weight  are  bases. 
All  pentavalent  hydroxids  are  acidic  in  nature,  but,  while  HN03  is  one 
of  the  strongest  acids,  HBi03  is  very  weak. 

NITROGEN.    N  — 14. 

Nitrogen  is  a  colorless,  tasteless,  odorless  gas,  existing  uncombiried 
in  the  atmosphere,  of  which  it  constitutes  about  four-fifths.  It  is  not 
appreciably  soluble  in  any  solvent.  One  of  its  chief  characteristics  is 
its  inertness.  It  requires  the  expenditure  of  considerable  energy  to  make 
it  unite  with  other  substances.  Although  its  true  character  is  much 
concealed  by  its  inactivity,  nitrogen  acts  as  a  typical  acid-forming  ele- 
ment. The  rare  hydronitric  acid,  HN3,  resembles  the  other  acid  hydrids. 
In  solution,  it  dissociates  into  the  ion  N3  to  a  much  greater  degree  than 
H2S  ionizes,  but  not  so  much  as  HF.  With  certain  metals,  nitrids  are 
formed — from  magnesium,  as  an  illustration,  Mg3N2 — which  are  decom- 
posed by  water,  liberating  NH3  and  forming  the  hydroxid  of  the  metal. 

Mg3N2  +  6HOH-*  3Mg(OH)2  +  2NH3 

These  compounds  are  of  scientific  importance  only.  United  with 
hydrogen  and  oxygen,  however,  nitrogen  forms  several  important  ions. 
The  detection  of  nitrogen  itself  comes  properly  under  the  subject  of 
gas  analysis  and  will  not  be  considered  here. 

Nitrogen  compounds  are  of  two  distinct  types.  In  one  the  element 
is  trivalent ;  in  the  other  it  shows  five  valences.  The  normal  tendency  is 
toward  the  formation  of  the  pentavalent  compounds,  trivalent  com- 
pounds being  generally  less  stable  than  the  higher  form  into  which  they 
tend  to  pass  quite  readily. 

With  hydrogen,  nitrogen  forms  compounds  of  a  basic  nature,*  the 
most  important  of  which  is  ammonia,  NH3.  It  is  a  colorless  gas  with  a 
pungent  odor,  reacting  readily  with  many  substances,  generally  by  direct 
addition,  forming  compounds  in  which  the  nitrogen  becomes  pentavalent. 
It  is  extremely  soluble,  one  litre  of  water  dissolving  nearly  800  volumes 
of  gas  at  ordinary  temperature  and  pressure,  forming  a  solution  which  is 

*  Except  hydronitric  acid,  HN8. 


NITKOGEN    COMPOUNDS.  147 

known  as  "ammonia  water."  Part  of  the  dissolved  NH3  unites  with 
the  water  to  form  a  compound,  NH4OH,  which  dissociates  into  OH  and 

+ 

NH4  ions  and  has  a  distinct  basic  reaction.  Ammonia  unites  with  es- 
pecial readiness  with  many  metallic  ions  forming  complex  ions  of 
various  types,  such  as  Ag.(NH8)2,  Cu.(NH3)4  and  Co.(NH3)6  (see  also 
cadmium,  chromium,  nickel,  zinc  and  mercury).  The  union  of  am- 
monia with  water,  or  ammonia  with  acids,  may  be  regarded  as  a  similar 
reaction,  i.  e.,  the  union  of  NII3  with  the  H  ion  to  form  the  NH4  ion,  as 
the  equations  show  : 


NH3  +  H20  -^NH4OH  NH3  +  HC1  -»  NH4C1 

NH3  +  H  +  OH-»NH4  +  OH  NH3  +  H  +  C1-»NH4  +  01 

NH3  +  H-»NH4  NH3  +  H->NH4 

Inasmuch  as  the  NH4  ion  is  very  stable  and  unites  with  the  anions 
of  acids  to  form  a  whole  series  of  salts  in  which  the  univalent  radicle, 
(NH4),  acts  as  a  metal,  the  radicle  is  called  ammonium,  the  ending 
1  '  ium  '  '  signifying  the  possession  of  metallic  properties.  Since  ammonium 
salts  most  closely  resemble  those  of  potassium,  the  former  are  treated 
in  detail  immediately  following  the  consideration  of  the  latter,  page  75. 

With  oxygen,  nitrogen  forms  five  compounds  all  of  which  are  more 
or  less  acidic  in  nature. 

Nitrogen  monoxid,  N20,  is  a  colorless  gas  with  a  sweet  taste  and 
odor,  somewhat  soluble  in  water.  Although  it  is  the  anhydrid  of  the 
rare  hyponitrous  acid,  (HNO)2,  it  does  not  unite  with  water  to  form  the 
acid.  At  ordinary  temperatures  the  gas  is  quite  stable,  neither  taking 
on  more  oxygen  nor  giving  up  its  own.  At  a  higher  temperature,  how- 
ever, it  decomposes  into  its  elements  and  becomes  a  good  oxidizing  agent. 

Nitrogen  dioxid,  N202  (as  it  is  often  called  and  written  for  the  sake 
of  systematization,  although  the  molecule  is  really  NO  and  contains  but 
one  atom  of  nitrogen  and  one  of  oxygen),  is  a  colorless  gas  not  appre- 
ciably soluble  in  water.  It  parts  with  its  oxygen  less  readily  than  any 
other  oxid  of  nitrogen  and,  therefore,  is  the  decomposition  product 
formed  most  frequently  when  higher  oxidized  compounds,  such  as  ni- 
trates and  nitrites,  break  down.  On  exposure  to  the  atmosphere,  it  readily 
takes  up  more  oxygen,  forming  brown  N204.  Nitrogen  dioxid  unites 
with  ferrous  salts  in  concentrated  solutions  to  form  a  very  characteris- 
tic dark  brown  compound,  with  ferrous  sulfate  having  the  composition 
(FeSO4)2.NO.  It  is  very  unstable,  however,  and,  on  warming  or  diluting, 
is  decomposed  and  the  color  disappears.  (See  pages  149  and  151.) 


148  QUALITATIVE    ANALYSIS. 

Nitrogen  trioxid,  N203,  may  be  obtained  as  a  blue  liquid  at  low  tem- 
peratures. At  zero  or  below,  it  decomposes  into  a  mixture  of  the  dioxid 
and  tetroxid,  hence  has  no  existence  at  ordinary  temperatures.  It  is  the 
anhydrid  of  nitrous  acid,  HN02. 

Nitrogen  tetroxid,  N204,  is  a  brown  gas  with  pungent  odor,  each 
molecule  of  which  is  entirely  dissociated  into  two  molecules  of  N02  at 
150°.  It  may  be  condensed  into  a  yellow  liquid,  boiling  at  26°.  It  is 
soluble  in  water,  reacting  with  it  to  form  a  mixture  of  nitrous  and  nitric 
acids  according  to  the  equation : 

N204  +  H20  ->  HN02  +  HN03 

It  is  a  strong  oxidizing  agent,  readily  giving  up  half  of  its  oxygen 
and  breaking  down  into  N2O2.  Since  the  dioxid  unites  with  the  oxygen 
of  the  atmosphere  to  reform  the  tetroxid  again,  these  substances  act  as 
"carriers"  of  oxygen  and  by  their  agency  many  compounds  may  be 
indirectly  oxidized  by  the  air  which  would  not  be  oxidized  directly  with- 
out them. 

Nitrogen  pentoxid,  N205,  is  a  colorless,  crystalline  solid,  melting  at 
30°  and  decomposing  at  50°  into  nascent  oxygen  and  N204.  It  is,  there- 
fore, a  fine  oxidizing  agent.  It  dissolves  in  water,  readily  uniting  with 
it  to  form  HN03.  It  is,  therefore,  nitric  acid  anhydrid  and  to  its  pres- 
ence the  oxidizing  action  of  nitric  acid  is  due. 

Nitrous  Acid,  HNO2. 

Nitrous  acid  is  a  compound  of  trivalent  nitrogen,  very  unstable 
in  character.  It  has  never  been  obtained  free  and  exists  to  a  slight 
concentration  only  in  aqueous  solution.  Acting  as  a  weak  monobasic 
acid,  it  forms  salts  which  are  quite  stable  both  in  and  out  of  solu- 
tion. On  acidifying  a  solution  of  a  nitrite,  the  free  acid  decomposes  for 
the  most  part  into  water  and  the  anhydrid,  N203,  the  latter  breaking 
down  further,  as  previously  stated.  That  part  of  the  nitrous  acid  which 
does  not  decompose  at  once  gradually  changes  into  nitric  acid  according 
to  the  equation : 

0 

// 
3H— 0— N=0    ->    H— O— N        +N202  +  H2O 

\ 
O 

When  heated  gently,  nitrites  are  unchanged  but,  when  ignited,  they 
break  down  into  the  oxids  of  nitrogen  and  the  metal. 


NITRITES.  149 

All  nitrites  are  readily  soluble  with  the  exception  of  the  silver  salt 
and,  on  boiling  with  Na2C03,  they  are  transposed. 

Nitrous  acid,  being  so  unstable,  yields  readily  to  chemical  agents  of 
either  an  oxidizing  or  reducing  nature.  With  KMn04  and  reagents  of 
a  similar  nature,  it  is  oxidized  to  nitric  acid,  each  molecule  of  nitrous 
acid  absorbing  one  atom  of  oxygen.  On  the  other  hand,  strong  reducing 
agents  are  oxidized  by  it,  two  molecules  of  nitrous  acid  giving  up  one 
atom  of  oxygen  and  breaking  down  into  N202  and  water  (see  pages  43 
and  47).  Thus  nitrous  acid  acts  either  as  an  oxidizing  or  reducing  agent. 

The  formation  of  nitrogen  dioxid  when  nitrous  acid  acts  as  an 
oxidizing  agent  gives  rise  to  the  dark  brown  compound  already  referred 
to  whenever  nitrous  acid  and  ferrous  salts  are  brought  together.  As  the 
nitrous  acid  oxidizes  part  of  the  ferrous  iron  to  the  ferric  condition,  the 
nitrogen  dioxid,  formed  during  the  reaction,  unites  with  the  unoxidized 
ferrous  salt  according  to  the  equations: 

2FeS04  +  2HNO2  +  H2S04  -»  Fe2(S04)3  +  N202  +  2H2O 
2FeS04  +  NO  -»  (FeSOJ  2.NO 

This  and  the  following  reactions  are  used  as  tests  for  nitrous  acid. 

Nitrous  acid  reacts  with  many  organic  compounds  to  form  highly 
colored  substances,  some  of  which  are  used  as  dyes.  With  phenylenedia- 
min  in  acid  solutions,  the  well-known  "Bismark  Brown"  dye  is  obtained, 
even  a  trace  of  nitrous  acid  giving  a  distinctly  yellow-colored  solution. 

2C6H4(NH2)2  +  HN02-^NH2C6H4N2C6H3(NH2)2  +  2H20 

With  sulfanilic  acid  and  naphtylamin  a  complex  reaction  of  a  similar 
nature  takes  place,  resulting  in  the  formation  of  compounds  of  an  in- 
tense red  color.  With  this  reagent  slight  traces  of  nitrous  acid  produce 
a  pink  coloration. 

The  nitrite  ion,  N02,  is  alone  of  importance.  Some  complex  ions 
are  formed  of  which  the  cobalti-nitrite  ion,  Co(N02)6  (see  page  231)  is 
the  best  known. 

The  barium  ion,  Ba,  gives  no  precipitate  with  the  nitrite  ion,  N02. 

The  silver  ion,  Ag,  gives  a  white  crystalline  precipitate  only  in  con- 
centrated solutions,  for  AgNO2  is  soluble  in  300  parts  of  H20. 

Nitric  Acid,  HNO3. 

Nitric  acid  is  a  compound  containing  pentavalent  nitrogen  correspon- 
ding to  nitrous  acid,  which  contains  trivalent  nitrogen.  Like  the  latter, 


150  QUALITATIVE    ANALYSIS. 

it  is  monobasic  but,  because  of  the  higher  state  of  the  oxidation  of  the 
nitrogen  it  contains,  its  acid  properties  are  much  stronger.  It  is  also 
much  more  stable  and  can  be  isolated  from  solution  as  a  colorless,  oily, 
fuming  liquid  boiling  at  86°,  with  partial  decomposition.  It  is  miscible 
with  water  in  all  proportions,  the  boiling-point  rising  steadily  as  the 
proportion  of  water  increases  until  at  120.5°  a  solution  containing  68 
per  cent,  of  HN03  distills.  This  solution  is  ordinary  "  concentrated 
nitric  acid."  If  the  solution  is  diluted  beyond  this  point,  the  boiling 
point  falls  as  the  proportion  of  acid  decreases  until  100°  is  reached,  the 
boiling-point  of  pure  water,  i.  e.,  of  0  per  cent.  HN03.  Since  all  mixtures 
of  nitric  acid  and  water  other  than  this  68  per  cent,  solution  boil  lower 
than  120.5°,  if  a  solution  of  nitric  acid  is  evaporated,  it  will  tend  to 
assume  this  definite  composition,  either  by  boiling  off  water,  if  it  is  less 
concentrated  than  68  per  cent.,  or  HN03  if  it  is  more  concentrated. 

"Concentrated  nitric  acid,"  on  standing  in  the  light,  decomposes 
slowly  into  water,  oxygen  and  nitrogen  tetroxid  which  colors  the  solution 
yellow  or  brown.  In  dilute  solution  it  is  stable,  however. 

The  salts  of  nitric  acid  are  naturally  less  readily  decomposed  than 
the  free  acid.  The  most  stable  nitrates  fuse  without  great  decomposition. 
When  strongly  ignited,  however,  all  decompose  with  the  liberation  of 
oxygen,  becoming  nitrites  first,  which  later  decompose  into  oxygen  and 
oxids  of  the  metal  and  of  nitrogen.  For  this  reason  nitrates  are  much 
used  as  oxidizing  agents  for  high-temperature  reactions. 


2K—  0—  N        +  H20  ->  2KOH  +  N202  +  30 

V 
o 

Heated  on  charcoal,  all  nitrates  deflagrate,  i.  e.,  the  carbon  is  oxidized 
rapidly  by  the  decomposing  nitrate  producing  flashes  of  light  and  slight 
explosions. 

All  normal  nitrates  are  soluble.  Certain  salts  of  nitric  acid,  such  as 
bismuth  and,  to  a  less  extent,  mercury,  react  with  water,  i.  e.,  hydrolyze 
and  form  basic  compounds  which  are  quite  insoluble.  All  nitrates  are 
transposed  by  boiling  with  sodium  carbonate. 

Pure  nitric  acid  is  one  of  the  best  oxidizing  agents  known.  The  "con- 
centrated" acid,  also,  possesses  strong  powers  of  oxidation.  "Fuming" 
nitric  acid  is  made  by  dissolving  N2O4  in  the  ordinary  concentrated  acid. 
It  possesses  the  additional  oxidizing  power  of  the  tetroxid  dissolved  in  it. 


NITEIC    ACID.  151 

On  diluting  the  acid  with  water,  its  oxidizing  action  decreases  as  its  ion- 
ization  increases.  Concentrated  nitric  acid,  for  instance,  oxidizes  hydri- 
odic  acid  very  readily  while  dilute  acid  separates  iodin  only  on  long 
standing.  In  other  words,  as  the  concentration  of  undissociated  acid 
falls,  the  oxidizing  action  decreases,  indicating  that  the  nitrate  ion,  NO3, 
is  much  more  stable  than  the  undissociated  acid,  to  which  we  must  ascribe 
the  greater  part  of  its  oxidizing  properties. 

To  what  degree  nitric  acid  will  part  with  its  oxygen  depends  very 
greatly  on  temperature,  concentration  and  the  nature  of  the  reducing 
substance.  By  abstracting  various  amounts  of  oxygen  all  the  oxids  of 
nitrogen  can  be  theoretically  formed,  the  element  obtained  and  hydrogen 
finally  added.  All  of  these  products  can  be  formed  practically.  Strong 
reducing  agents,  such  as  nascent  hydrogen,  under  suitable  conditions 
carry  the  reduction  all  the  way  down  to  ammonia.  Milder  reducing 
agents  will  not  carry  it  so  far.  Since  the  dioxid  is  the  stablest  of  the 
nitrogen  oxids,  under  ordinary  conditions  this  is  the  usual  reduction 
product  and  two  molecules  of  nitric  acid  liberate  three  atoms  of  oxygen. 
(See  pages  39  and  43.) 

Since  nitrogen  dioxid  is  formed  from  both  nitric  acid  and  nitrous 
acid,  the  test  with  ferrous  salts  already  outlined  for  nitrous  acid  (see 
page  149)  may  be  used  to  detect  nitric  acid  as  well,  the  reactions  being: 

6FeS04  +  2HN03  +  3H2S04  -»  3Fe2(SO4)3  +  N202  +  4H20 
2FeS04  +  NO  ->  (FeSOJ  2.NO 

Brucin  dissolved  in  concentrated  sulfuric  acid  is  colored  a  deep-red 
by  nitric  acid  (and  chloric  acid).  The  color  changes  rapidly  to  yellow 
and  finally  may  become  green. 

The  nitrate  ion,  N03,  is  the  only  known  anion,  for  the  reactions  of 
which  the  alkali  salts  serve  best. 

The  barium  ion,  Ba,  produces  no  precipitate. 

The  silver  ion,  Ag,  produces  no  precipitate. 

PHOSPHORUS.     P  — 31. 

The  element  exists  in  two  well-known  modifications.*  Ordinary 
phosphorus  is  a  very  pale-yellow  substance,  almost  transparent  and  of 
a  hard  waxy  consistency  at  ordinary  temperatures.  It  melts  at  44.4°  to 
a  yellow  liquid,  which  boils  at  290°.  In  carbon  disulfid  it  is  very  soluble 
but  in  other  solvents  it  dissolves  only  sparingly,  if  at  all.  Yellow  phos- 
phorus is  very  reactive,  uniting  directly  with  all  elements  except  nitro- 

*  Two  other  less  common  forms,  a  white  and  a  black  modification,  are  reported. 


152  QUALITATIVE    ANALYSIS. 

gen,  carbon  and  the  inert  argon  group.  With  the  metals  it  forms  phos- 
phids  which,  on  treatment  with  water  or  dilute  acids,  liberate  PH3.  With 
oxygen  it  unites  with  especial  readiness  and  it  is  necessary  to  keep  it  under 
water  to  protect  it  from  the  action  of  the  atmosphere.  Even  under  these 
conditions  it  becomes  covered  with  an  opaque  white  coating.  In  the  air  it 
glows  with  a  pale  light,  plainly  visible  in  a  darkened  room,  very  slight 
traces  being  sufficient  to  produce  a  marked  "phosphorescence."  At 
about  50°  phosphorus  is  spontaneously  inflammable,  burning  to  the  oxid, 
P205.  Heated  out  of  contact  with  the  air  to  a  temperature  near  its 
boiling-point,  it  is  converted  into  the  red  modification. 

Red  phosphorus  is  a  dark  granular  powder  which,  being  hygroscopic, 
tends  to  collect  together  in  balls  or  become  pasty.  It  does  not  melt  but, 
when  heated,  sublimes,  forming  a  vapor  which  condenses  into  liquid  yel- 
low phosphorus.  It  is  not  appreciably  soluble  in  any  solvent  and  at  ordi- 
nary temperatures  it  is  quite  inert.  By  heat  and  other  agents  it  may  be 
forced  to  react  and  is  then  converted  into  the  same  compounds  as  the 
other  modification. 

Either  form  of  phosphorus  may  be  detected  by  burning  and  collecting 
the  product  of  combustion  in  water,  which  may  then  be  tested  for  phos- 
phoric acid.  Nitric  acid  will  accomplish  the  same  change  but,  with  the 
yellow  modification,  concentrated  acid  acts  explosively. 

The  compounds  of  phosphorus  are  those  of  an  acid-forming  element. 
The  formation  of  halids  and  sulfids  may  be  considered  as  indicating  the 
possession  of  some  metallic  properties,  but  these  bodies  are  entirely  de- 
composed by  water  (hydrolyzed),  forming,  besides  hydrohalogen  or  hy- 
drosulfuric  acids,  hydroxids  of  phosphorus  which  are  typical  acids. 


3H20-»3HC1  +  P(OH)3   (or  H3P03) 

(or  H3P04  +  H2O) 


Phosphorus  compounds,  like  those  of  nitrogen,  are  of  two  general 
types,  trivalent  and  pentavalent,  the  former  tending  to  oxidize  readily 
into  the  latter.  For  this  reason  compounds  of  trivalent  phosphorus  are 
not  of  importance  in  analytical  chemistry. 

With  hydrogen,  several  compounds  are  formed  of  which  phosphin, 
PH3,  is  the  most  important.  Its  reactions  are  very  similar  to  those  of 
NH3  except  that  it  is  far  more  readily  oxidized  and  is  less  basic.  This 
point  is  indicated  by  the  fact  that  PH3  forms  salts  only  with  the  strong 
halogen  acids  and  these  salts  are  decomposed  by  water  with  the  libera- 
tion of  phosphin.  Phosphonium  hydroxid,  corresponding  to  NH4OH,  is 
unknown. 


PHOSPHORUS.  153 

With  oxygen,  phosphorus  forms  P203  and  P205  beside  the  less  im- 
portant P40  and  P204. 

Phosphorus  trioxid,  P203,  stands  in  the  same  relation  to  the  pentoxid 
as  does  N203.  Like  the  latter,  it  is  an  acid  anhydrid,  phosphorus  acid, 
H3P03  or,  minus  water,  HP02,  oxidizing  readily  to  the  higher  form. 
Phosphorus  trioxid  exists  to  a  large  extent  as  the  double  molecule  of  the 
form  (P2O3)2.  (See  also  arsenic  and  antimony,  pages  156  and  165.) 

Phosphorus  pentoxid,  P205,  is  a  light  snow-white  powder,  uniting 
with  water  with  extreme  readiness,  forming  phosphoric  acid.  Because 
of  this  property,  phosphoric  anhydrid  is  the  best  drying  agent  known  to 
science. 

The  simple  anion,  P,  is  little  known.  United  with  oxygen,  several 
ions  are  formed  of  which  those  given  by  phosphoric  acid  are  the  most 

important. 

Phosphoric  Acid. 

Phosphorus  pentoxid  reacts  with  water  and,  on  long  standing,  may 
form  phosphorus  pentahydroxid  in  solution  according  to  the  equation : 

P205  +  5H20->2P(OH)5 

Such  a  compound  cannot  be  separated  from  solution,  however.  At- 
tempts to  isolate  it  result  in  a  compound  containing  one  molecule  less 
water,  which  is  known  as  ortho-phosphoric  acid. 

OH 
/OH  H— 0 

//  \ 

P_OH    -»    H— 0— P  — O     +     H2O 

\\  / 

\OH  H— 0 

OH 

On  gentle  heating,  255°,  ortho-phosphoric  acid  loses  water,  forming 

pyro-phosphoric  acid. 

H— O 


+  H20 


154  QUALITATIVE    ANALYSIS. 

Pyro-phosphoric  acid  on  further  heating  breaks  down  into  meta- 
phosphoric  acid. 


->  211— 0— P        +  II20 


H— 0 


These  modifications  of  phosphoric  acid  are  more  stable  than  the  sim- 
ilar modifications  of  arsenic  (see  page  163)  or  other  acids  and  have 
an  independent  existence  for  a  time  in,  as  well  as  out  of,  solution. 

Meta-phosphoric  acid,  HP03,  is  a  colorless  glassy-looking  substance, 
which  is  often  cast  into  sticks  and  known  as  "glacial"  phosphoric  acid. 
It  is  readily  soluble,  in  aqueous  solution  acting  as  a  monobasic  acid,  of 
which  only  the  salts  of  the  alkaline  metals  are  readily  soluble.  It  may 
be  distinguished  from  the  ortho  and  pyro  acid  by  the  facts  that  it  gives 
white  precipitates  with  BaCl2  and  AgN03  solutions  and  will  coagulate 
egg  albumen. 

Pyro-phosphoric  acid,  H4P207,  resembles  the  meta  acid  in  appear- 
ance but  is  much  less  common.  It  may  be  distinguished  from  the  other 
forms  of  phosphoric  acid  by  the  fact  that  with  AgN03  it  gives  a  white 
precipitate,  Ag4P207,  only  in  solutions  which  are  nearly  neutral.  It 
does  not  coagulate  white  of  egg  and  gives  no  precipitate  with  BaCl2, 
except  in  an  alkaline  solution,  when  white  Ba2P207  falls. 

Both  of  these  partially-dehydrated  acids,  on  long  standing  in  aqueous 
solution,  take  up  water  slowly  and  form  the  ortho  acid.  The  reaction  is 
hastened  by  boiling  and  by  the  presence  of  a  little  dilute  HNO3,  under 
which  conditions  they  may  be  readily  changed  into  the  ortho  form. 
For  this  reason  they  will  not  be  considered  further. 

Ortho-phosphoric  acid,  H3P04,  is  a  colorless,  crystalline  solid,  which 
is  very  hygroscopic  and  extremely  soluble  in  water.  It  melts  at  38.6°. 
It  acts  as  a  tribasic  acid,  forming  salts  of  three  types,  primary,  secon- 
dary and  tertiary,  which  resemble  the  corresponding  salts  of  arsenic  acid 
(see  page  163).  All  the  alkaline  phosphates  (see  page  66)  are  soluble. 
The  primary  salts  of  the  alkaline  earths  and  of  magnesium  are  also  soluble. 


PHOSPHATES.  155 

Other  phosphates  are  generally  insoluble  in  water,  but  readily  soluble  in 
the  stronger  acids.  In  acetic  acid,  which  is  of  approximately  the  same 
strength  as  phosphoric  acid,  some  phosphates  are  soluble,  others  insoluble. 
Phosphoric  acid  (and  an  alkaline  phosphate  similarly)  in  solution 
dissociates  in  three  steps  as  indicated: 

H3P04->H  +  H2J04  (1) 

H2P04->H  +  HP04  (2) 

HP04-»H  +  PO4  (3) 

+ 
The  separation  of  H  ions  according  to  the  first  step  takes  place 

sufficiently  to  make  phosphoric  acid  an  acid  of  fair  strength.  lonization 
according  to  the  second  and  third  steps  is  so  slight,  however,  that,  were 
its  strength  dependent  upon  this  dissociation  primarily,  phosphoric  acid 
would  have  to  be  classed  among  the  very  weak  acids. 

In  a  solution  of  any  phosphate,  there  will  be  a  smaller  concentration 
of  the  simple  ion,  P04,  than  of  the  other  ions  formed  according  to  the 
first  or  second  steps.  Nevertheless,  when  metallic  ions  are  introduced  into 
a  solution  containing  all  three  of  these  phosphate  ions,  if  a  precipitate 
falls,  it  is  generally  of  the  type  of  normal  salt  formed  by  the  simple 
P04  ion  even  though  it  is  present  in  much  smaller  concentration.  For 
acid  salts,  such  as  would  be  formed  by  the  ions,  H2P04  and  HPO4,  are 
generally  much  more  soluble,  hence  not  so  readily  precipitated.  The 
secondary  salts  of  a  few  metals  are  quite  insoluble  and  are  precipitated 
on  the  addition  of  the  phosphate  ion.  Magnesium,  calcium,  strontium 
and  barium  (see  pages  93,  97,  97,  100)  are  of  this  class  and  precipitate 
normal  phosphates  of  the  type  Ba3(PO4)2  only  from  alkaline  solutions. 

The  barium  ion,  Ba,  precipitates  white  BaHP04  from  neutral  solu- 
tions, soluble  in  acids,  even  in  acetic  acid.  From  alkaline  solutions  only, 
Ba3(P04)2  is  precipitated.  This  salt  is  also  soluble  in  acetic  acid. 

The  silver  ion,  Ag,  precipitates  yellow  Ag3P04,  soluble  in  acids  and 
in  ammonium  hydroxid. 

These  two  reactions,  together  with  the  fact  that  egg  albumen  is  not 
coagulated,  serve  to  differentiate  ortho-phosphoric  acid  from  the  pyro  and 
meta  acids. 

The  magnesium  ion,  Mg,  precipitates,  from  solutions  containing  am- 
monium as  well  as  phosphate  ions,  white  floccy  (NH4)MgP04  which,  on 
boiling,  becomes  crystalline.  It  is  readily  soluble  even  in  acetic  acid. 
(See  page  93.) 

The  ferric  ion,  Fe,  precipitates  straw-colored  FeP04  which,  though 


156  QUALITATIVE    ANALYSIS. 

readily  soluble  in  the  strong  mineral  acids,  is  insoluble  in  acetic  and  other 

-t-H-  +++ 

organic  acids.  The  Al  and  Cr  ions  give  precipitates  with  similar  prop- 
erties. Phosphoric  acid  is  often  removed  from  solution  by  precipitating 
with  ferric  salts.  (See  pages  264  and  265.) 

The  molybdate  ion,  Mo04,  does  not  unite  directly  with  the  phosphate 
ion.  However,  when  a  solution  of  ammonium  molybdate,  (NH4)2MoO4, 
acidified  with  HNO3,  is  added  to  a  solution  containing  phosphoric  acid, 
the  latter  combines  with  the  anhydrid  of  molybdic  acid,  Mo03,  to  form 
several  complex  phospho-molybdic  acids,  with  compositions  ranging  from 
H3P04.10Mo03  to  H3P04.14Mo03.  Although  these  acids  are  soluble, 
their  ammonium  salts  are  insoluble  in  HN03,  a  reagent  added  to  prevent 
the  precipitation  of  any  phosphate  which  might  be  insoluble  in  water  or 
weaker  acids.  Under  appropriate  conditions,  a  yellow  precipitate  of 
the  proportions  (NH4)3P04.12Mo03  can  be  obtained,  the  formula  of 
which  may  be  taken  as  representing  the  composition  of  the  precipitate 
usually  brought  down.  These  complex  phospho-molybdates  are  broken 
down  in  alkaline  solutions  into  simple  phosphates  and  molybdates  of  the 
alkali  metals,  all  of  which  are  readily  soluble. 

Similar  complex  bodies  are  formed  by  arsenic  acid  with  Mo03.  (See 
page  164.) 

ARSENIC.     As  — 75. 

Arsenic,  as  it  ordinarily  occurs,  is  a  crystalline  gray  substance  with 
a  metallic  lustre.  Under  atmospheric  pressure,  it  does  not  melt  but  sub- 
limes, the  vapor  having  a  very  characteristic  garlic  odor.  At  ordinary 
temperatures,  it  is  but  little  acted  on  by  the  atmosphere  but  at  180° 
it  burns  readily  with  a  bluish-white  flame,  forming  the  white  trioxid, 
As2O3,  the  "  white  arsenic  "  of  commerce.  At  ordinary  temperatures, 
the  molecule  seems  to  double  up  and  exists  in  the  form  (As203)2. 

The  element  is  but  little  attacked  by  acids  other  than  HN03.  In  dilute 
solutions  of  this  acid,  arsenic  is  soluble.  Although  no  salts  of  acids  con- 
taining oxygen  are  definitely  known,*  As(N03)3  probably  exists  in  solu- 
tion to  some  extent.  Because  of  hydrolysis,  however,  by  far  the  greater 
part  of  the  arsenic  is  present  in  the  form  of  the  hydroxid. 

/N03      H— 0— H  /OH 

As— NO3  +  H— O— H  -*  3H— NO3  +  As— OH 

\N03      H— O— H  \OH 

Concentrated  HNO3  acts  as  an  oxidizing  agent  and  oxidizes  the  tri- 
valent  arsenious  compounds  to  pentavalent  condition  which  are  known 
*  Except  the  snlfate. 


AKSENIC.  157 

as  arsenic  compounds.  Possibly  As(N03)5  is  first  formed  but,  as  the 
metallic  properties  of  any  element  decrease  as  the  valence  rises  (see  page 
45),  this  salt  is  hydrolyzed  to  an  extent  even  greater  than  As(N03)3, 
and  practically  all  the  arsenic  exists  in  solution  in  the  form  of  the  pen- 
tavalent  hydroxid.  This  compound  loses  one  molecule  of  water,  as 
shown  in  the  equation,  and  exists  in  the  partially  dehydrated  form : 

/N03      H— O— H  /OH 

//NO,      H— 0— H  //OH  /OH 

As— N03  +  H— 0— H  ->  511— N03  +  As— OH  ->  0=As— OH  +  H20 
\\N03      H— 0— H  \\OH  \OH 

\N03      H— 0— H  \OH 

While  trivalent  arsenic  shows  the  properties  of  a  weak  base-forming 
element,  pentavalent  arsenic  exhibits  almost  no  basic  qualities  whatso- 
ever, but  does  possess  acid-forming  properties  quite  strongly  developed. 
Of  the  arsenic  ions,  the  trivalent  kathion,  As,  exists  only  to  a  small 
extent  in  solutions  because  of  the  great  degree  to  which  hydrolysis  takes 

H-  I  I  I  I 

place.  The  pentavalent  kathion,  As,  is  practically  unknown.  Arsenic 
enters  into  the  composition  of  several  anions,  of  which  those  derived  from 
arsenous  and  arsenic  acids  are  the  best  known.  These  will  be  considered 
separately  under  appropriate  headings. 

Compounds  of  arsenic,  of  whatever  nature,  are  readily  reduced. 
Pentavalent  compounds  are  rendered  trivalent  by  mild  reducing  agents, 
such  as  H2S  and  SO2.  More  active  reducing  agents  liberate  the  element 
itself,  while  the  strongest  reagents  carry  the  reduction  a  step  further 
and  add  hydrogen  to  it.  (See  page  39.) 

Heated  before  the  blowpipe  with  Na2C03  on  charcoal,  all  compounds 
of  arsenic  are  reduced  to  the  metal,  indicated  by  the  presence  of  the 
characteristic  garlic  odor.  The  hot  vapors  oxidize  as  they  leave  the 
heated  charcoal,  however,  and  usually  form  a  white  coating  of  the  oxid 
on  cool  objects. 

Heated  in  a  closed-tube  with  a  mixture  of  4  parts  Na2C03  to  1 
part  of  KCN,  all  arsenic  compounds  are  reduced  to  the  element  and  a 
ring  of  black  arsenic  will  sublime  out  of  the  mixture  and  condense  on  the 
cooler  portions  of  the  tube. 

In  the  "wet  way,"  SnCl2  causes  a  reduction  to  the  element  and  precip- 
itates arsenic  as  a  dark-brown  powder.  Under  the  name  of  the  Betten- 
dorf  Test,  this  reaction  serves  as  a  method  of  detecting  arsenic.  The 
reaction  takes  place  only  in  the  presence  of  strong  HC1,  however,  hence 
an  equal  volume  of  concentrated  HC1  is  added  to  the  solution  to  be  tested 


158  QUALITATIVE    ANALYSIS. 

before  adding  the  SnCl2  which,  for  this  purpose,  is  dissolved  in  concen- 
trated HC1  also.  In  dilute  HCI  solutions,  the  reduction  will  not  take 
place. 

Nascent  hydrogen  will  reduce  arsenic  compounds  to  the  lowest  stage, 
i.  e.,  it  will  form  arsenic  hydrici,  AsH3,  or  arsin.  This  is  a  colorless,  ill- 
smelling,  exceedingly  poisonous  gas,  analogous  to  NH3  in  composition 
but  far  less  stable.  Heated  alone,  it  decomposes  into  arsenic  and  hydro- 
gen. In  the  air,  ordinarily,  it  burns  to  H2O  and  As203.  If,  however, 
the  amount  of  air  is  limited,  it  burns  to  H2O  and  As. 

The  Marsh  Test,  the  most  delicate  method  known  for  detecting  com- 
pounds of  arsenic,  depends  on  these  reactions.  It  may  be  performed 
very  satisfactorily  as  follows:  Fit  a  six-inch  test-tube  with  a  stopper 
pierced  with  a  short  glass  delivery- tube  drawn  out  into  a  jet.  In  the 
bottom  of  the  test-tube  place  a  few  pieces  of  zinc*  and  cover  with  about 
10  ccm.  of  4N  H2S04  made  by  mixing  1  volume  of  concentrated  H2S04 
with  8  volumes  of  water.  Place  a  little  cotton  in  the  mouth  of  the  test- 
tube  to  hold  back  any  moisture,  insert  the  stopper  loosely  and,  after 
allowing  a  moment's  time  for  the  hydrogen  to  replace  all  air  in  the  test- 
tube,  cautiously  bring  the  jet  to  a  name  and  ignite  the  issuing  gas.  If  a 
slight  explosion  occurs,  it  indicates  that  all  of  the  air  had  not  been  driven 
out  of  the  test-tube  previously  to  igniting.  Replace  the  cotton  and  stop- 
per and  wait  a  longer  time  for  the  hydrogen  to  replace  the  air  completely, 
then  bring  to  the  flame  as  before.  If  the  gas  burns  quietly,  extinguish 
the  flame,  remove  the  stopper  and  cotton  plug  and  introduce  a  few  drops 
of  the  solution  to  be  tested.  Quickly  replace  the  cotton  and  stop- 
per and,  after  a  moment,  ignite  the  gas  as  it  issues  from  the  jet.f 

If  compounds  of  arsenic  be  present,  arsin  will  be  formed  and  will 
mix  with  the  hydrogen  and  burn  at  the  jet,  giving  the  flame  a  bluish- 
white  color  and  producing  faint  white  clouds  of  As2O3.  If  a  cold  porce- 
lain dish  be  held  against  the  flame,  thus  limiting  the  supply  of  air,  part 
of  the  arsenic  will  be  deposited  in  the  form  of  a  velvety  black  stain  which 
is  very  volatile  and  is  readily  soluble  in  NaOCl  (difference  from  Sb;  see 
page  166),  sodium  arsenate  being  formed. 

*  Since  As  is  very  widely  distributed  in  small  amounts,  many  laboratory  reagents, 
ordinary  zinc  especially,  contain  traces  of  it.  In  delicate  work  it  is  therefore  neces- 
sary to  make  a  "blank  test"  to  ascertain  whether  the  reagents  to  be  employed  are 
free  from  As.  The  method  of  procedure  is  exactly  as  described  except  that  the  solu- 
tion to  be  tested  is  not  added  to  the  hydrogen  generator.  If  no  stain  or  mirror  ap- 
pears, the  reagents  contain  no  As,  and  anything  which  may  develop  later  will  be 
entirely  due  to  the  solution  which  is  being  tested. 

t  Or  the  delivery-tube  may  be  made  longer  and  the  arsin  decomposed  by  heating 
the  tube  in  the  middle  portion.  Under  these  conditions,  the  As  is  deposited  as  a 
brilliant  mirror  on  the  sides  of  the  tube  at  some  distance  beyond  the  flame. 


AESENIC    SALTS.  159 

Metallic  copper  precipitates  arsenic  from  solutions  of  its  compounds 
strongly  acidified  with  HC1,  especially  on  warming.  The  arsenic  unites 
with  the  copper,  forming  copper  arsenid,  Cu5As2,  which  is  gray  in  color. 
This  is  known  as  the  Reinsch  Test.  Since,  however,  silver,  mercury,  bis- 
muth, gold,  platinum,  some  sulfur  compounds  and  certain  organic  sub- 
stances may  produce  similar  stains  or  deposits  on  copper,  such  an  indica- 
tion of  arsenic  must  be  confirmed  by  other  tests. 

The  Arsenious  Kathion,  As. 

The  hydroxyl  ion,  OH,  added  to  a  solution  containing  the  As  ion, 
produces  no  precipitate,  as  As(OH)3  is  fairly  soluble.  It  is  known  only 
in  solution,  however,  since  all  attempts  to  isolate  it  have  resulted  only  in 
the  dehydrated  product,  (As203)2.  This  is  soluble  in  water  to  a  slight 
extent,  forming  the  hydroxid  again,  but  it  dissolves  in  acids  or  alkalies 
more  readily,  since  the  salts  formed  by  arsenic  with  either  acids  or  bases 
are  more  soluble  than  the  hydroxid.  Arsenious  hydroxid,  As03H3, 

-H-+ 

may  ionize  in  two  ways,  either  as  a  base  into  the  ions  As  and  OH,  or  as 
an  acid  into  the  ions  As03  and  H.  Its  acidic  nature  is  somewhat  more 
strongly  developed  than  the  basic  and  its  solutions  in  pure  water  have_a 
slight  acid  reaction.  In  alkaline  solutions  the  high  concentration  of  OH 
ions  causes  a  repression  of  its  dissociation  as  a  base,  while  in  acid  solu- 

+ 

tions  the  high  concentration  of  H  ions  causes  a  repression  of  its  ioniza- 
tion  as  an  acid.  (See  page  31.) 

Consequently,  in  acid  solutions,  it  acts  as  a  base  while  in  alkaline 

+ 

solutions  it  behaves  as  an  acid.  In  acid  solutions,  the  H  ions  unite  with 
the  OH  ions  separated  by  As03H3  to  form  undissociated  water  and 
cause  a  continual  reduction  of  the  concentration  of  the  OH  ions  neces- 
sary to  maintain  equilibrium.  Hence  As03H3  continually  dissociates 
into  As  and  OH  ions  and  thus  is  forced  to  show  a  basic  character  (see 
page  36) .  The  properties  and  reactions  of  As03H3,  acting  as  an  acid,  will 

be  considered  later. 

===  +++ 

The  carbonate  ion,  C03,  has  no  effect  on  the  As  ion,  since  arsenic 

carbonate  is  unknown.  Alkaline  carbonate  solutions  are  always  hydro- 
lyzed  and  contain  the  OH  ion.  For  this  reason,  carbonate  solutions  have 
the  same  action  as  alkaline  hydroxids,  but  in  a  smaller  degree. 

The  sulfid  ion,  S,  gives  a  precipitate  of  yellow  As2S3  from  acid  solu- 
tions only;  from  alkaline  solutions  no  precipitate  is  thrown  down  be- 
cause of  the  formation  of  soluble  salts  of  thioacids  (see  page  162). 
Solutions  of  As203,  to  which  neither  acid  nor  alkali  have  been  added, 
color  slightly  when  H2S  is  introduced,  owing  to  the  formation  of  a  col- 


160  QUALITATIVE    ANALYSIS. 

loidal  precipitate.  They  remain  perfectly  clear  and  transparent,  how- 
ever, and  will  pass  through  the  finest  filter  paper.  On  boiling,  or  on 
the  addition  of  a  strong  electrolyte,  As2S3  is  precipitated  in  the  usual 

form.     (See  page  37.) 

+ 

The  H  ion  has  no  effect  on  As2S3,  but  HN03  oxidizes  it  to  H3As04. 
In  alkalies  and  their  sulfids  and  in  ammonium  carbonate,  As2S3  is  solu- 
ble, forming  salts  of  thioacids.  (See  page  162.) 

The  sulfate  ion,  S04,  produces  no  precipitate,  as  the  sulfate  is  soluble, 
especially  in  the  slightly  acid  solutions  necessary  to  prevent  the  forma- 
tion of  insoluble  basic  salts.  Other  salts  of  oxygen  acids  are  unknown. 

The  chlorid  ion,  01,  and  other  halid  ions  produce  no  precipitate, 
since  the  halid  salts  of  arsenic  are  soluble.  Arsenious  chlorid,  AsCl3  is 
an  oily  liquid,  B.  P.  130°.  Dissolved  in  a  little  water,  it  hydrolyzes  into 
the  basic  salt, 

OH 

/ 
As— OH 

\ 

Cl 

In  dilute  solutions,  it  separates  into  the  hydroxid  and  HC1  almost  en- 
tirely, owing  to  the  mass  action  of  the  water.  If  HC1  be  added  to  a 
dilute  solution,  however,  because  of  the  increased  concentration,  AsCl3 
is  formed  again  in  proportion  to  the  excess  of  acid  added.  An  equilibrium 
exists  which  may  be  represented  by  the  equation : 

2AsCl3  +  3H20  <-»  2As(OH)3  +  6HC1 

Because  of  these  and  similar  reactions,  if  any  arsenious  compound 
be  heated  with  HC1,  it  will  be  changed  into  AsCl3  which,  being  volatile, 
will  be  gradually  driven  off. 

Arsenic  Compounds. 

Pentavalent  arsenic  compounds  are  all  of  an  acid  nature.  No  salts 
of  oxygen  acids  are  known.  The  halids  are  entirely  decomposed  by 
water.  The  sulfid,  As2S5,  is  not  decomposed  by  water,  but  may  be 
precipitated  from  acidified  solutions  of  pentavalent  arsenic  compounds 
by  H2S.  It  is  yellow  in  color,  less  stable  than  the  trisulfid,  into  which  is 
slowly  decomposes  with  the  separation  of  sulfur.  It  is  insoluble  in  acids 
generally,  but  is  oxidized  by  HNO3  and  other  oxidizing  agents  to  H3As04 
and  H2SO4  (or  free  sulfur,  see  page  50).  It  dissolves  in  alkalies,  alkaline 


ARSENITES.  161 

sulfids  and  ammonium  carbonate,  forming  salts  of  thioacids  analogous 
to  those  of  trivalent  arsenic.     (See  page  164.) 

Arsenius  Acid. 

Arsenious  hydroxid,  As(OH)3,  shows  more  acidic  properties  than 
basic  properties,  and  is  ordinarily  known  as  arsenious  acid  and  written 
H3As03.  As  already  stated,  it  exists  only  in  solution  and,  on  any  attempt 
to  isolate  it,  separates  water  and  gives  the  oxid  which  thus  becomes  ar- 
senious acid  anhydrid.  But  slightly  soluble  in  water,  (As203)2  dissolves 
readily  in  alkaline  solutions,  forming  salts  of  which  three  types  are 
known.  They  are  derived  from  the  hypothetical  arsenious  acids  repre- 
sented by  the  formulae : 

0— H 

/ 
0— H  O  As— 0— H 

/  /  \ 

As— 0— H  As— O— H  0 

\  / 

0— H  As— 0— H 

O— H 

Ortho-arsenious  acid.          Meta-arsenious  acid.  Pyro-arsenious  acid. 

These  acids  differ  only  in  the  extent  to  which  they  are  hydrated. 
All  are  very  weak  and,  even  in  dilute  solution,  but  few  arsenious  anions 
are  found. 

Ortho-arsenites  are  not  very  stable,  Ag3As03  being  the  only  well- 
known  example.  Salts  formed  from  solutions  are  generally  of  the  meta- 
type,  MAs02,  although  some  pyro-arsenites  are  definitely  known. 

All  arsenites,  with  the  exception  of  salts  of  the  alkalies,  are  soluble 
with  difficulty  in  water,  but,  like  salts  of  weak  acids,  generally  are  readily 
soluble  in  the  strong  mineral  acids. 

All  arsenites  are  transposed  by  boiling  with  Na2C03  according  to 
the  equation  for  a  typical  reaction : 

2Ag3As03  +  3Na,C03  -»3Ag2C03  +  2Na3As03 

The  barium  ion,  Ba,  added  to  a  solution  containing  an  arsenite,  pre- 
cipitates white  barium  meta-arsenite,  Ba(AsO2)2,  readily  soluble  in 
dilute  acids. 

The  silver  ion,  Ag,  precipitates  yellow  (difference  from  H3As04,  see 
page  164)  silver  ortho-arsenite,  like  all  salts  of  arsenious  acid  soluble  in 

12 


162  QUALITATIVE    ANALYSIS. 

dilute  acids,  like  most  silver  salts  soluble  in  NH4OH  (see  page  85). 
Hence  it  requires  a  distinctly  neutral  solution  for  its  precipitation. 

The  magnesium  ion,  Mg,  in  the  presence  of  ammonium  salts,  gives  no 
precipitate  (difference  from  H3As04). 

The  molybdate  ion,  Mo04,  produces  no  precipitate  (difference  from 
H3As04). 

The  salts  of  thioarsenious  acid  are  very  analogous  to  the  salts  of 
arsenious  acid  and  may  be  considered  as  derived  from  the  latter  by  the 
substitution  of  sulfur  for  oxygen,  as  shown  by  the  formulae  : 


0—  NH4  S—  NH4 

/  / 

As—  0—  NH4  As—  S—  NH4 

\  \ 

0—  NH4  S—  NH4 

Ammonium  arsenite.  Ammonium  thioarsenite. 

Salts  of  this  nature  are  formed  when  arsenious  sulfid  is  dissolved  in 
alkalies,  alkaline  sulfids  or  in  ammonium  carbonate,  and,  also,  when  H2S 
is  passed  into  an  alkaline  solution  of  arsenious  acid,  as  is  shown  by  the 
following  reactions  : 

As2S3  +  6KOH  ->  K3AsS3  +  K3As03  +  3H20 

As2S3  +  3(NH4)2S*->2(NH4)3AsS3 

As2S3  +  3  (NHJ  2C03  -»  (NH4)  3AsS3  +  (NH4)  3As03  +  3C02 

2Na3As03  +  3H2S  ->  Na3AsS3  +  Na3As03  +  3H20 

On  acidifying  a  solution  of  a  thioarsenite,  thioarsenious  acid,  H3AsS3, 
is  first  formed  but,  just  as  arsenious  acid  splits  out  H20  and  forms  As203, 
so  thioarsenious  acid  splits  out  H2S  and  forms  As2S3  : 

2H3As03  -^  As203  +  3H20 

2H3AsS3  -»  As2S3  +  3H2S 

Consequently,  on  acidifying  a  solution  of  a  thioarsenite,  As2S3  is  pre- 
cipitated. Because  of  the  possible  formation  of  soluble  salts  of  thio- 
arsenious acid,  it  is  necessary  to  acidify  strongly  with  HC1  a  solution  of 
an  arsenite,  if  it  is  desired  that  As2S3  be  precipitated  by  introducing 
H2S. 

*  If  a  yellow  polysulfid  be  used  instead  of  a  colorless  monosulfid,  the  excess  of 
sulfur  will  oxidize  the  thioarsenite  over  to  the  corresponding  thioarsenate. 

3(NH4)2S.xS->2(NH4)2AsS4+  (3x  —  2)8 


ARSENATES.  163 

Arsenic  Acid. 

Pentavalent  arsenic  compounds,  as  previously  stated,  have  a  well- 
developed  acidic  character,  the  pentavalent  hydroxid  being  ordi- 
narily known  as  arsenic  acid.  It  is  most  readily  formed  by  the  oxi- 
dation of  arsenious  oxid  in  the  presence  of  water.  On  evaporation, 
ortho-arsenic  acid,  H3As04,  is  obtained  as  a  white,  crystalline  powder, 
readily  losing  water  on  heating,  giving  H4As207,  pyro-arsenic  acid,  and 
HAs03,  meta-arsenic  acid.  On  dissolving  in  water,  the  dehydrated  acids 
take  up  water  again. 

The  salts  of  ortho-arsenic  acid  are  of  three  types, — primary,  secon- 
dary and  tertiary — depending  on  the  amount  of  hydrogen  replaced  by 
metal.  Primary  salts  on  heating  form  meta-arsenates,  and  secondary 
salts  pyro-arsenates. 

M— 0  0 

\  / 

H— 0— As^O    ->    M— 0— As  +    H20 

/  \ 

H— 0  0 

M— 0 

M— 0— As=0  M— 0  0 

/  \     / 

H— 0  M— 0— As— 0     +    H,0 

->  / 

H— 0  M— 0— As 

\  /     \ 

M— 0— As=0  M— 0  0 

/ 
M— 0 

Solution  in  water  causes  hydration,  as  in  the  case  of  the  acids; 
hence,  the  acid  and  its  salts  are  always  of  the  ortho-  form  in  solution. 

Arsenic  acid,  although  dissociating  to  a  greater  degree  than  arsenious 
acid,  cannot  be  considered  strong,  and  its  alkali  salts  show  hydrolytic 
phenomena  to  a  greater  extent  than  the  corresponding  salts  of  phosphoric 
acid,  to  which  it  is  very  similar. 

Of  the  salts  of  arsenic  acid,  those  formed  by  the  alkali  metals  are 
readily  soluble.  Others  are  generally  insoluble  in  water,  but  readily 
soluble  in  mineral  acids. 

All  arsenates  are  transposed  by  boiling  with  sodium  carbonate. 


164  QUALITATIVE    ANALYSIS. 

The  barium  ion,  Ba,  added  to  a  solution  containing  an  arsenate,  pro- 
duces a  white  precipitate  of  Ba3(As04)2,  soluble  in  dilute  acids. 

The  silver  ion,  Ag,  precipitates  chocolate-colored  (difference  from 
H3As03,  see  page  161)  Ag3As04,  very  soluble  in  dilute  acids  and  in 

NH4OH,  requiring  a  distinctly  neutral  solution  for  its  precipitation. 

++ 

The  magnesium  ion,  Mg,  from  solutions  containing  the  ammonium 
ion  as  well,  precipitates  white  (NH4)MgAs04.6H20,  which  may  be  ren- 
dered beautifully  crystalline  by  boiling  or  by  being  brought  down  slowly 
from  solution.  The  presence  of  NH4C1  aids  the  transformation  from 
the  flocculent  to  the  crystalline  form. 

The  molybdate  ion,  Mo04,  from  solutions  of  the  ammonium  or  other 
similar  salt,  causes  a  yellow  precipitate  to  fall  which  has  the  composition 
represented  by  the  formula,  (NH4)3As04.12Mo03.4H20.  This  is  to  be 
regarded  as  the  ammonium  salt  of  a  very  complex  arseno-molybdic  acid. 
It  is  insoluble  in  HN03,  but  in  alkaline  solutions  it  dissolves  readily. 

The  salts  of  the  thioarsenic  acids  are  very  analogous  to  the  ordinary 
arsenates  and  resemble  closely  the  salts  of  thioarsenious  acid.  Two  types 
are  recognized,  depending  on  the  proportion  in  wrhich  the  oxygen  of 
arsenates  has  been  substituted  by  sulfur.  In  one  form,  all  of  the  oxygen 
has  been  replaced  by  sulfur;  in  the  other,  only  one-quarter,  as  is  shown 
by  the  formulae  : 

NH4—  O  NH4—  S  NH4—  0 

\  \  \ 

NH4—  0—  As=0        NH4—  S—  As=S        NH4—  0—  As=S 


NH4—  0  NH4—  S  NH4—  0 

Ammonium  arsenate.  Ammonium  thioarsenates. 

Salts  of  thioarsenic  acid  are  formed  by  dissolving  arsenic  sulfid, 
As2Sfc,  in  alkalies,  alkaline  sulfids  or  ammonium  carbonate  ;  by  dissolving 
arsenious  sulfid,  As2S3,  in  alkaline  polysulfids  ;  or  by  passing  H2S  into  a 
solution  of  an  arsenate  which  has  not  been  acidified.  The  following  equa- 
tions indicate  the  reactions  taking  place  : 

As2S5  +  6KOH  ->  K3AsS4  +  K3As03S  +  3H20 

As2S5  +  3(NH4)2S->2(NH4)3AsS4 

As2S5  +  3(NH4)2C03-+  (NH4)3AsS4+  (NH4)3As03S  +  3<702 
As2S3  +.3(NHJ2S.xS~>2(NH4)8AsS4  +  (3x—  2)S 
2Na3As04  +  5H2S  ->  Na3AsS4  +  Na3  As03S  +  5H20 

On  acidifying  a  solution  of  a  thioarsenate,  thioarsenic  acid,  H3AsS4 

(or  H3AsO3S),  is  first  formed,  but  it  immediately  decomposes  into  As2S5 


ANTIMONY.  165 

and  H2S,  analogously  to  the  manner  in  which  arsenic  acid  may  be  broken 
down  into  As205  and  H2O. 

2H3As04  -»  As205  +  3H20 
2H3AsS4-»As2S5  +  3H28 

On  acidifying  a  solution  of  a  thioarsenate,  therefore,  As2S5  is  precipi- 
tated. These  reactions  make  evident  the  reason  why  arsenic  sulfid  is  best 
precipitated  from  solutions  quite  strongly  acid  with  HC1.  On  passing  H2S 
into  an  acidified  solution  of  an  arsenate,  the  precipitate  usually  obtained 
consists  of  As2S5,  As2S3  and  free  sulfur,  indicating  that  the  H2S  acts 
partly  as  a  reducing  agent  as  well  as  a  precipitant. 

ANTIMONY  (Stibium).     Sb— 120.2. 

Antimony  is  a  white,  crystalline,  very  brittle  metal.  It  does  not  tar- 
nish in  the  air,  but,  when  ^eated,  burns  with  a  greenish-white  flame  to  the 
white  trioxid,  Sb203,  which,  except  at  extremely  high  temperatures,  has 
the  double  molecule,  (Sb203)2,  like  the  corresponding  oxid  of  arsenic. 
The  trioxid  readily  absorbs  more  oxygen,  if  heated,  and  becomes  Sb204. 

The  metal  is  not  ordinarily  acted  upon  by  acids  with  any  readiness 
even  when  heated.  The  oxidizing  action  of  HN03  is  more  affective  than 
HC1  or  H2S04  and  results  in  the  formation  of  (Sb203)2  in  dilute,  and 
Sb205  in  concentrated  solutions.  Aqua  regia  is  the  only  satisfactory 
solvent,  however,  converting  the  metal  into  chlorids. 

Antimony  requires  the  expenditure  of  a  small  amount  of  energy  to 
make  the  metal  pass  into  the  ionic  condition,  hence  it  is  precipitated  from 
solutions  of  its  salts  by  several  other  metals,  especially  zinc,  which  shows 
a  greater  tendency  to  become  ionic.  The  element  shows  a  valence  of 
three  and  five,  and,  as  is  always  the  case,  compounds  showing  the  lower 
valence  act  more  like  metallic  compounds  than  those  of  higher  valence. 
Trivalent  antimony  shows  considerably  more  basic  properties  than  tri- 
valent  arsenic,  but  its  acid  properties  are  not  so  well  developed.  Pen- 
tavalent  antimony  acts  as  a  very  weak  base-forming  element,  no  salts  of 
oxygen  acids  being  known.  Its  acid  properties  are  not  so  striking  as 
those  of  pentavalent  arsenic.  Since  the  trivalent  compounds  pass  quite 
readily  into  the  higher  form  in  the  presence  of  oxidizing  agents,  they 
are  to  be  regarded  as  reducing  substances. 

Under  the  influence  of  reducing  agents,  antimony  compounds  behave 
just  as  do  the  compounds  of  arsenic. 

Heated  before  the  blowpipe  with  Na2C03  on  charcoal,  all  antimony 
compounds  are  reduced  to  the  metal,  leaving  a  bright  white  globule 


166  QUALITATIVE    ANALYSIS. 

which  is  exceedingly  brittle  (difference  from  Sn,  Pb,  and  Ag),  does  not 
dissolve  in  HNO3  (difference  from  Pb,  Ag  and  Bi)  but  is  changed  to  a 
white  powder  (like  Sn).  It  will  not  dissolve  appreciably  in  hot  concen- 
trated HC1  (distinction  from  Sn)  and  passes  into  solution  only  on  treat- 
ment with  aqua  regia.  On  diluting,  H2S  gives  an  orange  precipitate. 
Since  the  metal  oxidizes  readily,  a  coating  of  white  oxid  forms  on  the 
charcoal  near  to  the  melt.  The  fumes  are  odorless,  thus  differentiating 
them  from  arsenic. 

i  Nascent  hydrogen  acting  on  any  soluble  antimony  compound  in  acid 
solution  forms  stibin,  SbH3.  In  alkaline  solutions,  it  does  not  form  (dis- 
tinction from  As).  It  is  very  similar  to  AsH3,  and,  like  it,  is  decom- 
posed on  heating  in  a  glass  tube  into  hydrogen  and  the  metal.  Since 
stibin  is  less  stable  than  arsin  and  antimony  is  less  volatile  than  arsenic, 
the  deposit  will  appear  much  nearer  to  and  usually  on  both  sides  of  the 
flame.  Stibin  burns  with  a  greenish-white  flame  to  the  oxid  but,  in  a 
limited  supply  of  air,  deposits  on  cold  porcelain  held  in  the  flame  metallic 
antimony  which,  unlike  the  brown  velvety  stain  produced  by  arsenic,  is 
black  and  metallic  in  appearance  and  will  not  dissolve  in  NaOCl. 

All  the  salts  of  antimony  hydrolyze  very  decidedly  and  tend  to  pre- 
cipitate oxy-compounds  from  solutions.  For  this  reason,  no  simple  salts 
of  antimony  are  soluble  in  water  in  the  ordinary  sense,  but  the  halids  are 
quite  soluble  in  the  presence  of  a  small  per  cent,  of  acid. 

Of  the  antimony  ions,  Sb  is  best  known.  The  pentavalent  kathion  is 
unknown.  Antimony  is  a  constituent  of  several  anions,  of  which  Sb03 
and  Sb04  may  be  considered  typical.  Corresponding  to  these  are  the 
thio  ions  SbS3  and  SbS4.  The  antimonate  ions,  in  which  the  element 
shows  a  valence  of  five,  are  better  known  than  the  antimonite  ions.  Com- 
plex anions  are  well  known.  Of  these,  the  antimony-tartrate  ion, 
SbO(C4H406)  is  the  most  important. 

The  reactions  of  the  Sb  ion  may  be  most  satisfactorily  studied  by 
using  slightly  acid  solutions  of  the  chlorid. 

The  Antimonous  Ion,  Sb. 

The  hydroxyl  ion,  OH,  added  to  a  solution  containing  the  Sb  ion, 
precipitates  Sb(OH)3  which  loses  water  in  various  proportions, 
forming  finally  (Sb203)2.  All  of  these  dehydration  products  are 
capable  of  separating  both  H  and  OH  ions,  the  former  being  called 
forth  in  alkaline  solutions  especially,  the  latter  in  the  presence  of 
acids.  For  this  reason,  the  precipitate  is  somewhat  soluble  in  alkalies  as 
well  as  in  acids.  Solutions  in  pure  water  react  slightly  acid,  hence  the 


ANTIMONY    SALTS.  167 

hydroxid  is  sometimes  called  ortho-antimonous  acid.  The  partially- 
dehydrated  substances  correspond  to  the  pyro-  and  meta-  acids.  The 
latter  has  never  been  definitely  isolated,  however.  Of  the  salts  of  the 
various  antimonous  acids  but  little  is  definitely  known.  Certain  ones  of 
the  meta  type,  M  —  0  —  Sb  =  0,  exist  but,  as  none  are  important,  they 
will  not  be  further  considered. 

Antimonous-  hydroxid  or  oxid  is  readily  soluble  in  tartaric  acid,  form- 
ing a  complex  antimony-tartrate  ion,  SbO(C4H406),  many  compounds 
of  which  are  soluble. 

OH  H 

/  \  \ 

Sb—  OH     +     H—  (C4H406)     -»     0  =  Sb—  (C4H406)     +     H20 


Bodies  containing  the  univalent  radicle,  (0  =  Sb  —  ),  are  known  as 
antimonyl  compounds.  The  potassium  antimonyl-tartrate  is  the  most 
important  and  is  known  as  "tartar  emetic"  because  of  its  medicinal 
action.  At  ordinary  temperatures,  this  complex  salt  is  soluble  in  about 
8  parts  of  water.  The  antimonyl-tartrate  ion  is  quite  stable  in  the  pres- 
ence of  OH  ions.  Hydrogen  ions  tend  to  decompose  it  and,  in  acid  solu- 

+-H- 

tions,  the  reactions  of  the  Sb  ion  predominate.  On  adding  HC1,  the  oxy- 
chlorid  is  precipitated.  Nitric  or  sulfuric  acid  bring  down  the  hydroxid 
because  of  hydrolysis  of  the  sulfate  and  nitrate  first  formed.  From  acid 
"tartar  emetic"  solutions,  the  same  precipitate  falls  when  the  S  ions  is 
introduced  as  from  SbCl3  solutions.  Similar  antimonyl  ions  are  formed 
with  citric  and  other  organic  acids. 

The  carbonate  ion,  C03,  produces  the  same  precipitate  as  the  OH  ion 
for,  since  antimony  carbonate  is  unknown,  it  is  the  OH  ions  formed  by 
hydrolysis  which  are  the  active  agents  in  solution. 

The  sulfid  ion,  S,  precipitates  orange  Sb2S3,  soluble  in  alkalies  and  in 
alkaline  sulfids  and  polysulfids,  forming  compounds  similar  to  those  of 
arsenic  (see  page  160).  Unlike  As2S3,  the  precipitate  is  soluble  in  hot 
HC1  and  insoluble  in  (NH4)2C03.  On  fusion,  it  turns  black  and  crys- 
tallizes in  the  form  of  the  natural  sulfid,  stibnite. 

The  sulfate  ion,  S04,  produces  no  precipitate,  since  Sb2(S04)3  is 
soluble  in  dilute  acid  solutions. 

The  chlorid  ion,  Cl,  produces  no  precipitate,  since  SbCl3  and  other 
halids  are  soluble  in  dilute  acid  solutions.  Antimony  trichlorid  forms 
usually  a  colorless  crystalline  mass,  which,  because  of  its  softness  at  ordi- 
nary temperatures,  is  known  as  "butter  of  antimony."  When  treated 
with  about  twice  its  weight  of  water,  it  hydrolyzes  and  forms  SbOCl,  but, 


0 
168  QUALITATIVE    ANALYSIS. 

if  diluted,  the  mass  action  of  the  water  disturbs  this  equilibrium  and 
hydrolysis  takes  place  to  a  greater  degree,  resulting  in  an  admixture  of 
the  oxid  with  the  oxychlorid.  This  is  a  reversible  reaction,  however,  and 
the  precipitate  dissolves  as  the  concentration  of  HC1  increases,  and  forms 
again  on  dilution.  The  precipitate  is  soluble  in  tartaric  acid  and  alka- 
line tartrates,  forming  the  complex  antimonyl-tartrates.  It  is  also  some- 
what soluble  in  alkalies,  forming  salts  of  antimonous  acid. 

The  thioantimonites  correspond  to  the  salts  of  antimonous  acid,  as 
do  the  thioarsenites  to  the  arsenites.  They  are  in  every  way  analogous 
to  the  thioarsenites  in  properties,  reactions  and  methods  of  formation. 
The  ion  SbS3  is  unstable  in  acid  solution  and,  on  acidifying  a  solution 
of  a  thioantimonite,  H2S  and  Sb2S3  is  precipitated  as  in  the  case  of  thio- 
arsenites. (See  page  162.) 

Pentavalent  Antimonic  Compounds. 

Pentavalent  antimony  acts  much  less  as  a  base-forming  element  than 
when  it  is  in  the  trivalent  condition,  hence  no  salts  of  oxygen  acids  are 
known  and  all  salts  hydrolyze  to  even  a  greater  extent  than  with  anti- 
monous compounds.  A  higher  percentage  of  acid  is  required  to  keep 
them  in  solution.  Antimony  pentachlorid,  SbCl5,  is  a  yellow  fuming 
liquid,  soluble  in  moderately  concentrated  HC1,  but,  in  such  solution, 
exists 'practically  as  undissociated  salt.  On  treatment  with  H2S  in  the 
cold,  orange  Sb2S5  can  be  precipitated.  Under  ordinary  conditions,  the 
precipitate  consists  of  Sb2S5,  Sb2S3  and  free  sulfur,  thus  indicating  that  a 
partial  reduction  to  the  lower  condition  has  been  caused  by  the  H2S.  An- 
timony pentasulfid  is  soluble  in  concentrated  HC1,  in  alkalies  and  alkaline 
sulfids  and  polysulfids,  forming  thioantimonates.  In  the  ordinary  solution 
of  (NH4)2CO3,  when  heated,  it  is  somewhat  soluble.  In  the  cold  it  is 
insoluble,  however,  and  may  be  differentiated  in  this  way  from  the  sul- 
fids of  arsenic.  (See  page  160.) 

Antimonic  Acid. 

The  pentavalent  hydroxid  of  antimony,  Sb(OH)5,  does  not  exist  but 
various  dehydration  products  are  known.  Since  these  all  possess  well- 
defined  acid  properties,  they  are  ordinarily  called  antimonic  acids. 
Three  forms  are  recognized,  corresponding  exactly  with  the  arsenic  acids. 


ANTIMONATES.  169 


H— O  I  0 

\ 

H— 0— Sb  =  O  0  H— O— Sb 

\ 
H— 0  \  0 


Ortho-antimonic  acid.  Pyro-antimonic  acid.  Meta-antimonic  acid. 


Many  salts  of  these  acids  are  known,  generally  of  the  pyro-  and 
meta-types.  Since  the  antimonic  acids  are  all  weak,  the  salts  all  hydro- 
lyze  considerably  and  are  generally  of  but  little  importance.  The  acid 
sodium  pyroantimonate,  Na2H2Sb207,  is  of  interest  because  of  the  fact 
that  it  is  one  of  the  most  insoluble  simple  sodium  salts  known.  It  re- 
quires about  400  parts  of  water  to  dissolve  it.  The  corresponding  potas- 
sium salt  is  more  soluble  and  precipitates  the  sodium  salt  when  the  Na 
ion  is  introduced.  This  is  sometimes  used  as  a  test  for  sodium  salts. 

The  barium  ion,  Ba,  added  to  a  solution  containing  an  antimonate, 
precipitates  white  Ba(Sb03)2. 

The  silver  ion,  Ag,  precipitates  AgSb03,  readily  soluble  in  acids  and 
in  NH4OH. 

Thioantimonates  are  better  known  than  most  other  salts  of  thioacids. 
In  general,  they  correspond  closely  with  thioarsenates  (page  164)  in 
properties  and  methods  of  formation.  The  thioantimonate  ion,  SbS4,  is 
no  more  stable  in  acid  solutions  than  is  the  ion  AsS4,  and  the  acid  decom- 
poses into  H2S  and  Sb2S5,  just  as  the  thioacids  of  arsenic  do.  The  most 
important  compound  of  this  type  is  known  as  ' '  Schlippe  's  salt ' '  and  has 
the  composition  indicated  by  the  formula  Na3SbS4.9H20. 

BISMUTH.     Bi  — 208.5. 

Bismuth  is  a  very  brittle,  crystalline  metal  of  a  reddish-white  color. 
At  ordinary  temperatures  it  is  not  tarnished  by  any  constituent  of  the 
atmosphere,  but,  at  higher  temperatures,  it  unites  with  oxygen,  forming 
pale  yellow  bismuth  trioxid,  Bi203. 

The  metal  is  but  little  soluble  in  dilute  HC1  or  H2S04.  Dilute  HN03 
has  some  action.  In  hot  concentrated  H2S04  or  fairly  strong  HN03  it 
dissolves  readily.  The  element  absorbs  a  small  amount  of  energy  in 
passing  from  the  metallic  into  the  ionic  condition,  therefore  its  salts  may 


170  QUALITATIVE    ANALYSIS. 

be  reduced  without  difficulty.  The  metal  is  precipitated  from  solutions 
of  its  salts  by  metallic  zinc. 

Heated  before  the  blowpipe  with  Na2C03  on  charcoal,  all  bismuth 
salts  are  reduced,  leaving  a  bright  metallic  globule  which  is  brittle  (differ- 
ence from  Ag,  *Pb,  and  Sn).  A  coating — orange-yellow,  hot;  lemon- 
yellow,  cold  (difference  from  Pb  and  Sn) — is  formed  on  the  charcoal, 
due  to  oxidation  of  part  of  the  metal.  The  globule  is  soluble  in  warm, 
fairly  concentrated  HN03  and,  with  HC1,  gives  a  precipitate  (if  at  all) 
only  in  very  dilute  solutions  (difference  from  Ag  and  Pb).  This  precipi- 
tate is  soluble  in  a  slight  excess  of  acid  and  may  by  this  means  be 
further  differentiated  from  lead  and  silver. 

Salts  of  bismuth  are  generally  colorless.  None  of  them  dissolve  read- 
ily in  pure  water,  owing  to  the  formation  of  insoluble  oxysalts.  Bismuth 
acts  as  a  very  weak,  trivalent,  base-forming  element,  hence  its  salts 
hydrolyze  very  considerably.  The  regular  equation  would  be : 

N03  OH 

/  / 

Bi— N03  +  3HOH  -+  Bi— OH  +  3HN03 

\  \ 

N03  OH 

Hydrolysis  is,  however,  a  typical  reversible  reaction,  for  bismuth 
hydroxid  is  soluble  in  HN03  with  the  formation  of  the  nitrate  and  water. 
Hence  a  condition  of  equilibrium  ensues  in  which  both  bismuth  hydroxid 
and  nitrate  are  present.  The  mixture  seems  to  be  within  the  molecule 
(see  page  21),  however,  and  basic  bismuth  nitrate  results.  This  loses 
water,  as  the  equation  shows,  and  forms  the  oxysalt  which  is  known  as 
the"subnitrate." 

OH  00 

/  //          // 

Bi— OH    -*    Bi— 0— N         +    H2O 

\  \ 

N03  0 

To  prevent  the  precipitation  of  salts  of  this  character,  an  acid  must  be 
present  in  the  solution.  In  the  presence  of  a  little  of  their  respective 
acids,  the  chlorid,  sulfate  and  nitrate  all  dissolve  readily.  For  this  rea- 
son, all  bismuth  salt  solutions  react  strongly  acid  to  litmus. 

+-H- 

The  only  bismuth  ion  is  Bi.  This  shows  little  tendency  toward  the 
formation  of  complex  ions,  hence  bismuth  reactions  are  for  the  most  part 
normal.  The  nitrate,  chlorid  or  sulfate  are  satisfactory  solutions  for 

•H-+ 

studying  the  reactions  of  the  Bi  ion. 


BISMUTH    SALTS.  171 

The  hydroxyl  ion,  OH,  gives  a  flocculent  precipitate  of  white 
Bi(OH)3,  under  ordinary  conditions  insoluble  in  an  excess  of  alkaline 
hydroxids.  In  hot  concentrated  KOH  or  NaOH  it  is  slightly  soluble, 
however,  thus  showing  a  tendency  to  act  as  an  acid-forming  element.* 

On  boiling,  Bi(OH)3  loses  water  and  forms  Bi  <QTT  which  is  slightly 

yellowish  in  color.    On  further  heating,  Bi203  is  formed. 

When  heated  with  an  alkaline  solution  of  sodium  stannite,  Na2Sn02, 
the  oxid  or  hydroxid  turns  black,  owing  to  the  formation  of  a  mixture  of 
a  lower  oxid,  BiO,  and  of  metallic  bismuth,  both  of  which  are  black.  The 
tin  is  oxidized  as  the  bismuth  is  reduced,  according  to  a  reaction  which 
may  be  written  so  as  to  express  both  stages  of  the  reduction  : 

Bi0        2NaSnO  -»  BiO      Bi      2NaSn0 


23 


The  carbonate  ion,  C03,  brings  down  basic  bismuth  carbonate,  owing 
to  hydrolysis  of  the  precipitating  reagent. 

The  sulfid  ion,  S,  precipitates  brownish-black  Bi2S3,  soluble  in  HNO, 
due  to  oxidation  to  Bi(N03)3  and  free  sulfur,  but  otherwise  insoluble  in 
acids.  Unlike  antimony,  arsenic  and  tin,  the  acid  properties  of  bismuth 
are  not  sufficiently  developed  to  allow  the  formation  of  salts  of  this  acid, 
therefore,  Bi2S3  is  not  soluble  in  alkalies  or  alkaline  sulfids. 

The  sulfate  ion,  S04,  causes  no  precipitate,  since  Bi2(S04)3  is  soluble. 
An  oxysalt,  with  the  formula  given  below,  forms  readily  in  dilute 
solutions. 


\  \ 

S04  Cr04 

/  / 

Bi=0  Bi=0 

The  chromate  ion,  Cr04,  precipitates  a  yellow  oxychromate  (formula 
given  above),  soluble  in  HNO3  but  insoluble  in  NaOH  and  thereby  dif- 
ferentiated from  yellow  PbCr04  which  is  soluble. 

*  A  few  compounds  are  known  in  which  bismuth  is  pentavalent.  According  to 
the  general  rule,  as  the  valence  rises  the  acidic  properties  develop.  The  compound 


H  —  O  —  Bi^    ,  similar  to  H  —  O  —  N^     and  H  —  O  —  P<^    ,is  known  as  bis- 

muthic  acid.  It  is  of  a  bright  red  color  and  forms  salts,  among  which  an  oxid  Bi2O4 
(or  BiO2)  is  to  be  considered  as  the  oxy  bismuth  salt  of  bismuthic  acid,  as  its  struc- 
tural formula  indicates,  Bi  =  O  Bi  =  O 


172  QUALITATIVE    ANALYSIS. 

The  chlorid  ion,  01,  gives  no  precipitate,  since  BiCl3  is  soluble.    The 

,o 

tendency  toward  the  formation  of  bismuth  oxychlorid,  Bi  y     ,  is  very 

XC1 

marked  whenever  the  concentration  of  acid  becomes  sufficiently  low  to 
permit  hydrolysis  to  take  place.  A  beautiful  illustration  of  the  phe- 
nomena of  equilibrium  and  mass  action  may  be  obtained  by  diluting 
cautiously  a  BiCl3  solution  until  a  precipitate  forms.  This  will  dissolve 
on  the  addition  of  a  drop  or  two  of  concentrated  HC1,  to  be  precipitated 
again  as  soon  and  as  often  as  the  concentration  of  acid  falls  below  the 
equilibrium  requirement  because  of  dilution.  Since  bismuth  is  more  of 
a  base-forming  element  than  antimony,  the  precipitated  BiOCl  is  far  more 
soluble  in  acid  than  the  similar  salt,  SbOCl,  which  it  closely  resembles. 
Since  bismuth  is  less  of  an  acid-former  than  antimony,  BiOCl  is  insoluble 
in  alkalies,  while  SbOCl  dissolves  readily.  Bismuth  oxychlorid  is  insol- 
uble in  tartaric  acid,  H2(C4H406),  and  is  turned  black  by  H2S,  whereas 
SbOCl  is  soluble  in  tartaric  acid  and  is  turned  orange  by  H2S. 

The  iodid  ion,  I,  precipitates  black  BiI3  from  fairly  concentrated 
solutions.  From  dilute  solutions,  orange  BiOI  falls,  an  oxyiodid  similar 
to  BiOCl.  Bismuth  tri-iodid  is  soluble  in  excess  of  alkaline  iodids  forming 
the  complex  ion,  BiI4,  which  decomposes  on  dilution  into  the  simple  ions 
with  the  precipitation  of  black  BiI3  first  and  later  of  orange  BiOI. 

The  cyanid  ion,  ON,  probably  forms  Bi(CN)3  at  first,  but  this  is 
hydrolyzed  completely  and  Bi(OH)3  is  precipitated.  No  simple  cyanids 
of  trivalent  metals  are  known. 

The  phosphate  ion,  P04,  precipitates  white  granular  BiP04  which  is 
interesting  from  the  fact  that,  unlike  all  ordinary  phosphates,  it  is  not 
readily  soluble  in  solutions  of  the  strong  acids. 

Group  5,  Second  Type. 

Vanadium  51.2;  Niobium  94;  Tantalum  183. 

Three  rare  elements  are  the  only  known  members  of  this  type.  In 
general,  they  resemble  those  elements  of  the  first  type  which  have  medium 
atomic  weights  and  they  will  not  be  considered  further. 

Group  6,  First  Type. 
Oxygen  16;          Sulfur  32.1;          Selenium  79.2 ;          Tellurium  127.6. 

The  elements  of  Group  6  show  characteristic  valences  of  six  or  two. 
The  members  of  this  and  the  following  group,  however,  are  capable  of 


OXYGEN    GBOUP.  173 

forming  compounds  in  which  the  elements  exhibit  valences  intermediate 
between  those  which  are  considered  as  most  characteristic.  Inasmuch 
as  the  properties  of  the  compounds  change  as  the  valence  of  the  element 
rises,  compounds  of  widely  different  nature  may  be  formed  by  the  same 
element. 

The  members  of  this  group  are  light  non-metallic  elements,  not  par- 
ticularly active  at  ordinary  temperatures  but,  as  the  temperature  rises, 
they  unite  readily  with  most  other  elements.  For  these  reasons,  they 
are  found  in  nature  both  free  and  combined.  Hydrids,  of  the  type  H20, 
are  well  known  and  important.  They  possess  the  properties  of  very 
weak  acids.  Combination  with  oxygen  takes  place  readily  in  various 
proportions,  oxids  being  formed  which  act  as  acid  anhydrids.  These 
unite  with  water,  forming  acids,  the  strength  of  which  increases  as  the 
valence  of  the  element  in  them  rises,  those  of  the  sulfuric  acid  type, 
H2S04,  being  much  stronger  than  the  sulfurous  acid  type,  H2S03. 

Selenium  and  tellurium  are  not  of  common  occurrence  and  they  will 
not  be  considered  further. 

OXYGEN.     0  — 16. 

Oxygen  is  a  colorless,  odorless  gas,  occurring  free  in  the  atmosphere 
to  the  extent  of  about  21  per  cent,  by  volume.  In  combination  with 
other  elements,  it  is  estimated  that  oxygen  constitutes  about  50  per  cent, 
of  the  known  crust  of  the  earth.  It  is  but  slightly  soluble  in  water,  yet 
dissolves  in  this  medium  more  than  in  any  other  ordinary  solvent. 

Although  oxygen  enters  into  combination  with  every  element  except 
fluorin  and  the  members  of  the  inactive  argon  group  (see  page  65), 
at  ordinary  temperatures  it  is  comparatively  inactive  and  most  reactions 
take  place  so  slowly  as  to  be  almost  imperceptible.  At  higher  tempera- 
tures, however,  it  becomes  very  active  and  it  unites  with  many  substances 
so  rapidly  that  both  heat  and  light  are  produced.  Union  with  oxygen  is 
the  chemical  reaction  ordinarily  understood  by  the  term,  combustion. 

Oxygen  itself  does  not  burn  under  ordinary  conditions,  but  other 
substances  burn  in  it,  and  the  presence  of  oxygen  gives  to  the  atmosphere 
its  power  to  support  combustion.  As  might  be  expected,  all  substances 
burn  more  rapidly  in  pure  oxygen  than  in  air,  and  the  ability  to  hasten 
combustion  is  taken  as  a  test  for  the  presence  of  oxygen.*  A  piece  of 
charcoal,  glowing  in  the  air,  bursts  into  flame  when  introduced  into  a 
jar  containing  oxygen  and  many  substances,  such  as  iron,  which  do  not 
burn  in  air,  burn  vigorously  in  pure  oxygen. 

*  Some  of  the  oxids  of  nitrogen  support  combustion  better  than  air  but  they  may 
be  readily  differentiated  from  oxygen  by  their  odors. 


174  QUALITATIVE    ANALYSIS. 

An  allotropic  modification  of  the  element  is  known  as  ozone.  It  is  a 
gas  with  a  very  characteristic  penetrating  odor,  and  pale  blue  color  seen 
only  in  large  volumes  of  the  gas.  It  may  be  formed  by  subjecting  the 
ordinary  modification  to  the  action  of  electricity,  when  a  contraction  in 
volume  occurs.  In  general,  whenever  oxygen  is  formed,  some  ozone 
is  formed  also.  It  is  comparatively  unstable  at  ordinary  tempera- 
tures and,  on  heating  to  300°,  breaks  down  completely  into  the  ordinary 
modification  with  an  increase  in  volume  half  as  great  as  the  original  vol- 
ume. For  these  and  other  reasons,  the  molecule  of  ozone  is  supposed  to 
be  03  while  that  of  ordinary  oxygen  is  O2. 

Ozone  is  about  as  active  at  ordinary  temperatures  as  oxygen  is  at 
high  temperatures.  It  is  one  of  the  most  energetic  oxidizing  agents 
known,  usually  giving  up  only  one-third  of  its  oxygen  and  changing  into 
the  ordinary  form.  Moist  ozone  oxidizes  all  the  commoner  metals  except 
gold  and  platinum  to  the  highest  possible  valence  and  affects  many  of 
the  non-metallic  elements  in  the  same  way. 

Oxids  are  formed  by  the  direct  union  of  elements  with  oxygen  (or 
ozone)  and  also  by  heating  the  corresponding  hydroxids,  carbonates, 
nitrates,  or  salts  of  organic  acids  which  usually  decompose  into  carbonates 
first.  Many  oxids  are  insoluble  in  water.  Of  these,  oxids  formed  by 
metals  are  soluble  in  the  strong  acids,*  while  those  formed  by  non-metal- 
lic elements  are  soluble  in  solutions  of  the  strong  bases. 

All  oxids  unite  with  water  to  form  hydroxids,  either  directly  or 
indirectly. 

Na20  +  H20  -»  2NaOH  CaO  +  H20  ->  Ca  ( OH)  2 

PA  +  5H20-»  2P(OH)'  SO/+  3H20 -*  S(OH)6  * 

Although  the  properties  of  these  hydroxids  grade  one  into,  the  other 
very  gradually,  for  convenience  they  may  be  divided  into  two  general 
classes : 

(a)  Those  hydroxids  which  in  aqueous  solution  separate  H  ions  are 
known  as  acids,   and  the   elements  from  which  they  are   derived  as 
non-metallic  or  acid-forming  elements. 

(b)  Those  hydroxids  which  in  aqueous  solution  separate  OH  ions  are 
known  as  bases,  and  the  elements  from  which  they  are  derived  as  metals 
or  base-forming  elements. 

The  hydroxids  of  certain  elements  do  not  separate  H  or  OH  ions  to 

*  Certain  naturally-occurring  oxids,  or  the  artificial  products  which  have  been 
strongly  ignited,  are  soluble  with  extreme  difficulty. 


OXYGEN    COMPOUNDS.  175 

% 

any  great  extent  but,  to  a  slight  extent,  they  may  separate  either  under 
different  conditions,  i.  e.,  they  may  act  either  as  acids  or  as  bases,  depend- 
ing on  their  environments.  In  the  presence  of  more  strongly  acidic 
hydroxids,  they  act  as  bases  and  in  the  presence  of  distinctly  basic  hy- 
droxids  they  play  the  part  of  acids.  (See  page  22.) 

Many  elements  unite  with  oxygen  in  more  than  one  proportion  and 
the  chemical  nature  of  the  corresponding  hydroxids  varies  with  the  state 
of  oxidation,  the  more  highly  -oxidized  compounds  possessing  more  acidic 
properties  than  those  in  which  the  element  exhibits  a  smaller  valence 
(see  page  45).  As  the  valence  of  the  element  rises,  usually  the  full 
amount  of  water  necessary  to  form  the  normal  hydroxid  does  not  unite 
with  the  oxid  but  the  commonly-occurring  forms  may  be  represented  as 
the  normal  hydroxids  minus  one  or  more  molecules  of  water. 

P(OH)5  —  H2O^H3P04  S(OH)6  —  2H20-*H2S04 

The  different  degrees  of  hydration  are  denoted  by  the  prefixes  ortho, 
meta  and  pyro,  these  terms  being  used  most  commonly  in  connection  with 
acidic  hydroxids.  Ortho  acids  represent  the  ordinary  degree  of  hydra- 
tion, i.  e.,  the  form  most  commonly-occurring.  Meta  acids  are  formed 
from  ortho  acids  by  the  loss  of  a  molecule  of  water.  Pyro  acids  are 
obtained  by  the  separation  of  one  molecule  of  water  from  two  molecules 
of  the  ortho  acids. 

ORTHO  TYPE.  META  TYPE.        PYRO  TYPE. 

H—  0 

\ 
H—  0—  B         —  H2O  ->  H—  0—  B=  0 

H—  O 

H—  0 

H—  0\  H—  O 

Si         _  H0  ->  Si=0 


H—  O/  H—  0 

H—  0 

H—  0  O 

\  / 

H—  0—  P=  O  —  H,0  -»  H—  O—  P 

/  \ 

H—  O  O 


176  QUALITATIVE    ANALYSIS. 

OETHO.  PYRO. 


H— O  0 

\    / 

2  S        —  H20 

/    \ 

H— O  O 


H— 0 

2  H— 0— P=0       -  H20 
H— 0 


Water,  H2O. 

Water,  the  common  oxid  of  hydrogen,  is  a  clear,  almost  colorless, 
tasteless  liquid,  boiling  at  100°  and  changing  to  a  colorless  solid,  ice,  at 
0°.  It  dissolves  a  great  many  substances  readily  and  is  the  most  general 

solvent  known.    It  may  be  regarded  as  the  hydroxid  of  hydrogen  and  as 

+  — 

such  it  dissociates  to  an  exceedingly  small  degree  into  H  and  OH  ions.* 
It  is,  therefore,  either  acidic  or  basic  in  nature.  With  most  bodies,  it  acts 
as  a  neutral  reagent,  however,  but  with  substances  possessing  extreme 
properties,  water  takes  on  a  definite  character  of  its  own. 

In  contact  with  the  strong  base-forming  elements  water  acts  as  an 
acid.  These  metals  show  a  much  greater  tendency  to  pass  into  the 
ionic  condition  than  hydrogen  does.  When  sodium  or  potassium  is 
thrown  upon  water  the  electrical  charge  passes  from  the  H  ions  of  the 
water  to  the  metal,  which  then  enters  the  ionic  condition.  On  losing  its 
charge,  the  hydrogen  passes  out  of  the  ionic  condition  and  is  liberated  in 

*  There  is  evidence  to  indicate  that  the  dissociation  of  water  into  H  and  OH  ions 
is  only  the  first  step  in  the  process  and  that  water,  like  other  weak  acids,  undergoes 
further  dissociation  to  a  very  slight  degree  with  the  formation  of  O  ions  and  a  second 
H  ion. 


WATEK.  177 

gaseous  form.  A  comparison  of  the  equations  expressing  the  action  of 
water  and  of  hydrochloric  acid  on  metallic  sodium  will  show  that  water 
acts  as  a  typical  acid. 

2Na  +  2H  +  2  OH  -*  2Na  +  2  OH  +  H2 
2Na  +  2H  +  2C1  ->  2Na  +  2C1  +  H2 

Strong  acid-forming  elements,  on  the  other  hand,  cause  water  to  play 
the  part  of  a  base.  Fluorin  and  chlorin  show  a  much  greater  tendency 
to  assume  a  negative  charge  than  the  hydroxyl  radicle  does.  When  these 
gases  are  passed  into  water,  the  charge  passes  from  the  hydroxyl  ion  of 
the  water  to  the  halogen,  which  enters  the  ionic  condition.  The  hydroxyl 
radicles  pass  from  the  ionic  condition  and  decompose  into  water  and 
oxygen  as  the  equation  shows: 

4:F  +  4H  +  4  OH  ->  4H  +  4F  +  2H20  +  02 

Although  water  dissociates  to  an  extremely  small  degree  (only  one 
or  two  parts  in  a  thousand  million  being  ionized),  it  causes  substances 
dissolved  in  it  to  dissociate  to  a  greater  degree  than  any  other  common 
solvent  does  (see  page  13).  This  fact,  taken  in  connection  with  the 
great  solvent  power  of  water,  makes  aqueous  solution  a  most  important 
medium  for  the  investigation  of  chemical  reactions. 

Hydrogen  Dioxid,  Hydrogen  Peroxid,  H202. 

Hydrogen  dioxid  is  a  colorless  oily  liquid,  boiling  at  85°  under  70 
mm.  pressure.  When  heated  under  ordinary  atmospheric  pressure,  it 
decomposes  with  explosive  violence  into  water  and  oxygen.  The  liquid 
reacts  acid  to  litmus  and  is  miscible  with  water  in  all  proportions,  a  3 
per  cent,  solution  being  the  ordinary  hydrogen  peroxid  of  commerce.  In 
solution  it  is  much  more  stable  than  in  the  pure  state. 

Hydrogen  dioxid  stands  in  relation  to  water  about  as  ozone  does  to 
oxygen.  At  ordinary  temperatures,  its  solutions  decompose  slowly  into 
water  and  oxygen  and,  in  the  presence  of  reducing  agents,  this  decompo- 
sition is  frequently  rapid.  Therefore,  hydrogen  peroxid  is  a  good  oxi- 
dizing agent,  each  molecule  of  peroxid  giving  up  an  atom  of  oxygen. 

In  alkaline  solutions,  iron,  cobalt  and  manganese  are  oxidized  to  the 
trivalent  condition  and  chromium  passes  from  the  trivalent  base-forming 
condition  into  the  hexavalent  acid-forming  condition.  Cobaltic  oxid, 
Co203,  and  manganese  dioxid  decomposes  hydrogen  peroxid  very  rapidly 
and  cause  an  abundant  evolution  of  oxygen. 
13 


178  QUALITATIVE    ANALYSIS. 

In  acid  solutions,  chromium  is  oxidized  one  step  higher  than  in 
alkaline  solutions  and  becomes  the  blue  heptavalent  perchromic  acid  (see 
page  192).  This  reaction  is  frequently  used  as  a  test  for  hydrogen  dioxid 
although  this  reagent  is  less  delicate  and  less  reliable  than  titanium  sul- 
fate  for  the  detection  of  H202.  In  this  latter  test,  the  colorless  tetra- 
valent  salt  of  titanium  is  oxidized  to  the  hexavalent  condition,  which  is 
characterized  by  a  yellow  or  orange  color.  (See  page  306.) 

Although  a  typical  oxidizing  agent,  hydrogen  dioxid,  because  of  its 
instability,  reacts  with  certain  other  oxidizing  agents  in  such  a  way  that 
both  are  reduced  and  oxygen  gas  liberated.  For  this  reason,  perman- 
ganate is  bleached  according  to  the  equation: 

5H202  +  2HMn04  +  2H2S04  ->  2MnS04  +  8H20  +  502 

Similarly  hypochlorous  acid  is  reduced  to  hydrochloric  acid  and 
perchromic  acid  is  reduced  to  trivalent  chromium  by  an  excess  of 
hydrogen  peroxid  (see  page  267).  In  alkaline  solutions,  also,  copper, 
silver  and  gold  salts  are  reduced  to  the  metals  and  oxygen  is  evolved. 

The  salts  of  hydrogen  dioxid  are  the  metallic  dioxids  or  peroxids. 
Only  a  few  are  known.  Of  these,  the  peroxids  of  the  alkali  metals  only 
are  readily  soluble;  the  peroxids  of  the  alkaline  earths  are  much  less 
soluble  and  others  are  highly  insoluble. 

Sodium  peroxid,  Na202,  is  a  heavy  straw-colored  powder  which  unites 
with  water  to  form  a  hydrate,  Na202.8H20,  so  energetically  that  much 
heat  is  liberated.  It  is  hydrolyzed  considerably,  however,  so  that  its 
solution  in  water  is  equivalent  to  an  alkaline  solution  of  H202.  It  is  used 
as  a  very  strong  oxidizing  agent. 

Barium  dioxid,  Ba02,  treated  with  acids,  is  the  source  of  commercial 
hydrogen  peroxid. 

Ba02  +  H2S04  -»  BaS04  +  H2O2 

Manganese  dioxid,  Mn02,  and  lead  dioxid,  Pb02,  are  typical  higher 
oxids.  When  warmed  with  hydrochloric  acid,  their  equivalent  of  oxygen 
is  used  in  oxidizing  the  acid,  and  chlorin,  instead  of  oxygen,  is  liberated. 

Mn02  -(-  4HC1  — >  MnCl2  -j-  2H20  -f-   C12 

The  barium  ion,  Ba,  causes  no  precipitate,  for  Ba202,  although  quite 
insoluble,  is  not  formed  in  the  ' '  wet  way. ' ' 

The  silver  ion,  A+g,  causes  no  precipitate,  for  Ag202  is  not  formed  in 
the  ' '  wet  way. ' ' 


SULFUE. 

SULFUR.     S  —  32.1. 

The  element  sulfur  exists  in  several  modifications.  The  ordinary 
form  consists  of  a  light  yellow,  crystalline  substance  melting  at  114.5  a 
to  an  amber  fluid,  which  becomes  darker  and  more  viscous  as  the  tem- 
perature rises  to  250°.  On  higher  heating,  it  becomes  mobile  again  and 
boils  at  450°.  It  is  not  dissolved  appreciably  in  ordinary  solvents,  but  is 
readily  soluble  in  2  or  3  parts  of  carbon  disulfid.  Another  modification 
of  different  crystal  form,  also  soluble  in  CS2,  is  known.  It  is  stable  only 
above  95°.  At  least  two  amorphous  forms  are  recognized,  one  soluble 
in  CS2;  the  other,  "plastic  sulfur,"  is  elastic  and  insoluble  in  CS2. 
Sulfur  precipitated  by  the  decomposition  of  sulfur-containing  com- 
pounds is  generally  soluble  in  CS2.  "Flowers  of  sulfur"  is  a  finely 
divided  form  obtained  by  rapid  cooling  of  the  vapor.  It  is  partly  soluble 
and  partly  insoluble  in  CS2. 

All  modifications  of  sulfur  afford  for  the  most  part  the  same  chemical 
reactions.  They  dissolve  in  warm  solutions  of  alkaline  hydroxids,  sulfids 
and  sulfites,  entering  into  chemical  reactions  in  each  case,  according  to 
the  equations: 

6NaOH  +  4S  ->  2Na2S  +  Na2S203  +  3H20 

xNa2S  +  xS  -»  Na2S2,  Na2S3,  Na2S4,  Na2S5 

Na2S03  +  S  -»  Na2S203 

Sulfur  unites  directly  with  most  elements,  generally  requiring  tem- 
peratures higher  than  ordinary  for  appreciable  reaction. 

Sulfur  burns  with  a  pale  blue  flame  to  sulfur  dioxid,  S02,  a  colorless 
gas  (B.  P.  — 8°)  with  a  very  sharp,  pungent  odor,  used  as  a  disinfectant 
and  bleaching  agent.  It  is  the  anhydrid  of  sulfurous  acid.  More  than 
forty  volumes  of  the  gas  dissolve  in  one  volume  of  water,  all  of  which 
may  be  removed  again  on  boiling  the  solution.  It  oxidizes  slowly  in  the 
air  at  ordinary  temperatures,  more  rapidly  when  heated  or  in  the  pres- 
ence of  oxidizing  agents,  to  S03. 

Sulfur  trioxid  exists  in  two  different  forms.  One  modification  is  a 
liquid, — M.  P.  15° ;  B.  P.  46° — which  changes,  on  standing,  into  a 
white,  crystalline  solid  which  sublimes  without  melting.  As  produced  in 
the  laboratory  usually,  S03  appears  as  heavy  white  fumes  without  par- 
ticular odor  which  produce  violent  coughing  when  inhaled.  It  is  the 
anhydrid  of  sulfuric  acd. 

Free  sulfur  is  usually  detected  by  burning  it  to  S02,  which  may  be 
recognized  readily  by  its  odor,  or  the  oxidation  product  may  be  collected 
in  water  and  detected  by  any  of  the  tests  for  sulfurous  acid.  In  organic 


180  QUALITATIVE    ANALYSIS. 

bodies  which  do  not  ionize,  sulfur  is  detected  by  converting  it  into  com- 
pounds which  dissociate  into  some  one  of  the  well-known  ions  containing 
sulfur.  (See  pages  308  and  309.) 

The  simple  sulfur  ion,  S,  may  be  considered  as  the  anion  of  the  acid, 
H2S.  United  with  oxygen,  sulfur  forms  several  important  ions  which 
will  be  considered  under  their  corresponding  acids. 

Hydrosulfuric  Acid,  Hydrogen  Sulfid,  H2S. 

Hydrogen  sulfid  is  a  colorless  gas  with  a  very  characteristic  dis- 
agreeable odor  like  that  of  decaying  eggs.  It  acts  as  a  poison  on  the 
human  system,  dilute  chlorin  being  the  best  antidote.  It  burns  with  a 
pale  blue  flame  to  water  and  S02.  At  ordinary  temperatures,  it  oxidizes 
readily,  either  to  water  and  free  sulfur,  or  to  H2SO4  as  the  final  product. 
Therefore,  it  is  a  strong  reducing  agent,  each  molecule  of  H2S  absorbing 
either  one  or  four  atoms  of  oxygen.  It  is  fairly  soluble  in  water  (2  or 
3  volumes  of  gas  in  one  of  water)  ,  its  solution  being  known  in  the  labora- 
tory as  "  hydrogen  sulfid  water"  and,  when  occurring  in  nature,  as 
"sulfur  water."  By  boiling  the  solution,  all  H2S  may  be  removed. 

Hydrogen  sulfid  in  aqueous  solution  acts  as  a  very  weak  acid,  dis- 
sociating slightly  according  to  the  first  equation  and  to  an  extremely 
small  degree  according  to  the  second  equation. 


(1)     H2S-»H  +  HS  (2) 

Consequently  there  are  always  more  HS  than  S  ions  in  its  solutions,  but, 
since  acid  salts  are  much  more  soluble  than  normal  salts,  the  normal 
sulfids  are  precipitated  from  solutions.  The  sulfids  of  the  alkali  metals 
are  soluble  and  suffer  hydrolysis.  The  sulfids  of  the  akaline  earth  metals 
are  fairly  insoluble,  but,  in  the  presence  of  water,  are  changed  by 
hydrolysis  into  acid  sulfids  which  are  readily  soluble.  The  sulfids  of 
magnesium,  aluminum,  chromium  and  ferric  iron  are  soluble,  but,  owing 
to  the  fact  that  the  hydroxids  of  these  metals  are  insoluble,  the  sulfids  are 
completely  hydrolyzed  in  the  presence  of  water  and  the  hydroxids  are 
precipitated.  The  sulfids  of  cobalt,  nickel,*  manganese,  zinc  and  ferrous 
iron,  though  insoluble  in  water,  are  soluble  in  the  stronger  acids  and 
hence  are  not  precipitated  in  the  presence  of  HCL  The  sulfids  of  silver, 
mercury,  lead,  copper,  bismuth,  cadmium,  arsenic,  antimony,  tin,  gold 
and  platinum  are  precipitated  in  the  presence  of  dilute  HC1,  but  of  these 
the  sulfids  of  lead,  cadmium,  antimony  and  tin  are  soluble  in  more  con- 
*  For  the  peculiar  action  of  CoS  and  NiS  toward  HC1  see  pages  229  and  233. 


HYDROGEN    SULFID.  181 

centrated  acid,  while  the  others  do  not  dissolve  in  HC1  at  all,  contrary 
to  the  general  rule  that  salts  of  weaker  acids  dissolve  in  solutions  of 
stronger  acids  (see  page  35).  The  varying  action  of  H2S  on  salts  of  the 
various  metals  is  taken  advantage  of  in  qualitative  analysis  and  is  the 
basis  of  the  separation  into  groups.  (See  page  248.) 

The  differing  action  of  HC1,  or  the  H  ion  in  general,  on  these  different 
sulfids  depends  principally  on  the  relative  solubilities  of  these  sulfids  in 
water.  Those  which  are  highly  insoluble  in  water  are  not  soluble  in  HC1, 
while  those  which  are  most  soluble  in  water  dissolve  most  readily  in  HC1. 
To  illustrate :  When  H2S  is  passed  into  a  solution  of  a  soluble  copper  salt 
acidified  with  HC1,  there  are  present  Cu,  S  and  H  ions.  Two  distinct 
conditions  of  equilibrium  will  exist  in  the  solution  (see  page  24),  one  with 
regard  to  CuS,  another  with  regard  to  H2S.  In  each  there  will  be  one 
common  term,  the  S  ion.  Hydrogen  sulfid  is  quite  soluble  and  the  solu- 
tion may  be  saturated  with  the  gas  and  a  definite  equilibrium  obtained 

+ 

between  the  dissolved  but  undissociated  H2S  and  the  ions  H  and 
S.  Since  H2S  is  weakly  acidic,  the  concentration  of  the  ions  will 
be  represented  by  small  factors,  while  that  representing  the  undis- 
sociated substance  will  be  relatively  large.  Should  the  concentration 
of  H  and  S  ions  become  greater  than  is  necessary  for  the  stated  con- 
dition of  equilibrium,  undissociated  H2S  will  be  formed,  the  solution 
will  become  supersaturated  and  the  gas  will  bubble  out. 

On  the  other  hand,   copper  sulfid  is  extremely  insoluble   in   H20, 

++  = 

hence  only  very  small  concentrations  of  Cu  and  S  ions  can  exist  in  the 

solution  together.  There  is,  however,  a  high  concentration  of  Cu  due  to 
the  original  soluble  copper  salt.  When  H2S  is  introduced,  the  concentra- 
tion of  sulfid  ions  becomes  too  great  for  equilibrium.  Undissociated  CuS 
is  formed  in  such  quantities  that  the  solution  becomes  supersaturated  and 
precipitates  CuS  long  before  it  becomes  saturated  with  H2S. 

Conversely  CuS  does  not  dissolve  even  in  concentrated  HC1,  for  the 
concentration  of  S  ions  necessary  to  separate  CuS  from  solution  is  less 
than  that  required  to  separate  H2S.  Therefore  a  saturated  solution  of 
CuS  never  contains  S  ions  in  sufficient  quantity  to  separate  H2S. 

Cadmium  sulfid  is  considerably  more  soluble  than  CuS,  hence  requires 
a  higher  concentration  of  S  ions  to  precipitate  it  from  solution  than  does 
CuS.  With  the  increased  concentration  of  S  ions,  a  lower  concentration 
of  H  will  be  sufficient  to  cause  the  separation  of  H2S  from  solution. 
Nevertheless  the  concentration  of  H  ions  given  by  dilute  HC1  is  not  suf- 
ficiently great  to  cause  the  separation  of  H2S  rather  than  CdS.  There- 


182  QUALITATIVE    ANALYSIS. 

fore,  CdS  is  precipitated  from  solutions  of  cadmium  salts  by  H2S  in  the 
presence  of  not  more  than  4  per  cent,  of  HC1.  If  the  concentration  of 
HC1  be  greater  than  this,  CdS  is  not  precipitated.  On  the  other  hand, 
CdS  is  dissolved  readily  by  concentrated  HC1,  for  the  concentration  of  S 
ions  given  by  a  saturated  solution  of  CdS  is  greater  than  is  necessary  for 
equilibrium  with  regard  to  the  H2S.  Undissociated  H2S  is  formed,  the 
solution  becomes  supersaturated  and  the  gas  bubbles  out.  By  this  de- 
crease in  the  concentration  of  the  S  ion,  the  equilibrium  with  regard  to 
CdS  is  disturbed  and  more  CdS  goes  into  solution  to  restore  the  concen- 
tration of  S  ions.  This  causes  further  liberation  of  H2S  which,  in  turn, 
causes  further  solution  of  CdS. 

Zinc  sul-fid  is  more  soluble  than  CdS,  hence  requires  a  higher  concen- 
tration of  S  ions  to  cause  its  precipitation  from  solution.  But  with  the 

increased  concentration  of  S  ions,  a  correspondingly  smaller  concentra- 

+ 

tion  of  H  ions  will  be  sufficient  to  cause  the  separation  of  H2S.  The  con- 
ditions necessary  to  separate  ZnS  differ  but  little  from  those  necessary  to 
separate  H2S.  The  concentration  of  H  ions  given  by  HC1  is  sufficiently 
great  to  exceed  the  equilibrium  conditions  for  H2S  before  those  for  ZnS 
are  reached.  Hence,  ZnS  cannot  be  precipitated  in  the  presence  of  HC1. 
But,  on  the  contrary,  ZnS  dissolves  readily  in  dilute  HC1 ;  for  the  concen- 
tration of  S  ions  given  by  a  saturated  solution  of  ZnS  is  sufficiently  great 

to  cause  the  separation  of  H2S. 

+ 

The  concentration  of  H  ions  given  by  acetic  acid,  however,  is  much 
less  than  that  furnished  by  HCL  If  H2S  is  passed  into^a  zinc  solution 
slightly  acidified  with  acetic  acid,  the  concentration  of  S  ions  necessary 
to  precipitate  ZnS  will  be  reached  before  it  becomes  great  enough  to 
separate  H2S;  i.  e.,  before  the  solution  *  becomes  saturated  with  H2S. 
Therefore,  ZnS  is  precipitated  and,  once  thrown  down,  it  does  not  dis- 
solve in  acetic  acid. 

Ferrous  sulfid  is  still  more  soluble  than  ZnS  and  it  cannot  be  precipi- 
tated in  the  presence  of  acetic  acid,  for  a  saturated  solution  of  H2S  does 
not  contain  sufficient  S  ions  to  cause  the  separation  of  FeS.  The  alkaline 
sulfids  are  much  more  soluble  than  H2S,  however,  and  their  solutions 
afford  S  ions  in  sufficient  quantity  to  cause  the  separation  of  FeS. 

Conversely,  FeS  will  dissolve  in  acids  readily,  for  its  saturated  solu- 
tion contains  sufficient  S  ions  to  cause  the  separation  of  H2S.  In  the 
laboratory,  H2S  is  generally  prepared  by  the  action  of  acid  on  FeS,  this 
sulfid  being  used  because  of  its  cheapness. 

Sodium  sulfid  is  exceedingly  soluble.  Its  solutions  contain  very  high 
concentrations  of  S  ions,  so  great,  in  fact,  that  the  concentration  of  H 


SULFIDS.  183 

ions  due  to  dissociated  water  is  sufficiently  great  to  exceed  the  equilibrium 
for  H2S  and  cause  some  of  the  undissociated  substance  to  form,  the  odor 
of  which  may  be  detected  in  the  solution.  This  constitutes  the  hydrolysis 
of  soluble  sulfids. 

These  facts  may  be  summarized  and  tabulated  in  the  following  man- 
ner. The  condition  represented  is  that  equilibrium  for  H2S  already 
discussed,  in  which  any  increase  in  the  concentration  of  H  or  of  S  ions 
will  cause  the  separation  of  H2S  from  solution.  The  large  letters  repre- 
sent high  concentrations,  small  letters  correspondingly  low  concentrations. 

Concentration    Concentration 
Salt.  solubility. 


Na2S  \      I*     /'  H  Fromwater 

FeS            \     1     /  JsJ  H  From  acetic  acid 

ZnS              \  I  '  /  S  H  From  very  dil-  HC1 

CdS,             \W  S  JJ  From  4  per  cent.  HC1 

CuS  V         .      s  No  acid  gives  sufficient  H  ions 

Sulfids,  not  soluble  in  HC1,  may  be  decomposed  by  heating  with  con- 
centrated H2S04.  The  H2S  set  free  is  oxidized  by  the  hot  H2S04,  how- 
ever, into  water  and  free  sulfur  ;  and  SO2,  from  the  reduction  of  H2S04, 
is  given  off  instead  of  H2S. 

H—  0  OH  0 

\  //  \  // 

S         4-          S    ->     8         +S  +  2H20 

/  \  /  \ 

H—  0  OH  0 

Sulfur  may  be  detected  in  sulfids  of  this  character  by  fusing  with 
NaOH,  when  transposition  takes  place.  Much  of  the  Na2S  formed  will 
be  oxidized  during  the  fusion,  but  sufficient  will  remain  in  the  aqueous 
extract  to  give  reactions  for  the  sulfid  ion. 

On  ignition  out  of  contact  with  the  air,  most  sulfids  fuse  unchanged. 
When  heated  in  the  atmosphere,  the  sulfur  frequently  burns  out,  leaving 
an  oxid,  or  the  sulfid  may  be  oxidized  to  a  sulfate. 

All  sulfids  formed  by  precipitation  are  transposed  by  boiling  with 
Na2C03. 

The  barium  ion,  Ba,  gives  no  precipitate  with  solutions  containing 
sulfids  under  ordinary  conditions,  for,  although  BaS  is  somewhat  in- 


184  QUALITATIVE    ANALYSIS. 

soluble,  it  hydrolyzes  to  BaSH,  and  then  to  Ba(OH)2  and  H2S,  all  of 
which  are  soluble. 

The  silver  ion,  Ag,  precipitates  black  Ag2S,  insoluble  in  most  of  the 
ordinary  reagents.  Hot  HN03  oxidizes  it  to  the  soluble  Ag2S04  and  a 
large  excess  of  warm  KCN  will  dissolve  it,  forming  KAgC2N2  and  K2S. 

•f-4- 

The  lead  ion,  Pb,  precipitates  black  PbS,  soluble  in  concentrated  HC1 
and  in  dilute  HN03  on  warming.  A  filter-paper  moistened  with  lead 
acetate  is  often  used  to  detect  fumes  of  H2S,  but  sodium  plumbite, 
Na2Pb02,  is  a  more  delicate  reagent  for  the  same  purpose.  It  may  be 
prepared  by  adding  NaOH  or  KOH  to  a  soluble  lead  salt  until  the 
Pb(OH)2  first  precipitated  is  redissolved. 

Sodium  nitroferricyanid,  or  nitroprussid,  Na2(Fe  =  C5N5NO),  pro- 
duces a  transient  purple  color  in  alkaline  solutions  of  sulfids  but  not  in 
acid  or  neutral  solutions. 

Metallic  silver,  also  copper  and  mercury,  are  blackened  by  the  action 
of  H2S  or  alkaline  sulfids  in  the  presence  of  water  and  oxygen  (air) 
accorcjjng  to  the  equation : 

4Ag  +  2Na2S  +  02  +  2H20  ->  2Ag2S  +  4NaOH 

A  solution  of  H2S  in  water  becomes  turbid  on  standing,  owing  to  oxi- 
dation and  separation  of  free  sulfur,  as  has  been  stated  previously.  Un- 
der similar  conditions,  an  alkaline  sulfid  does  not  become  turbid,  although 
oxidation  takes  place  in  the  same  way.  The  sulfur  set  free  by  the  action 
of  oxygen  reacts  with  the  alkaline  sulfid  to  form  polysulfids,  of  which 
the  following  types  are  known,  Na2S2,  Na2S3,  Na2S4,  Na2S5.  The  acids 
corresponding  to  these  salts  are  very  unstable  and,  on  acidifying  a  solu- 
tion of  a  polysulfid,  H2S  is  liberated  and  free  sulfur  precipitated  for  the 
most  part.  Under  appropriate  conditions,  however,  a  yellow  oil  may  be 
obtained  which  consists  of  hydrogen  and  sulfur  in  the  proportions  corre- 
sponding to  a  mixture  of  the  hydrogen  polysulfids.  It  soon  decomposes 
into  H2S  and  sulfur.  The  alkaline  polysulfids  are  important  in  analysis. 
See  pages  252  and  253. 

Sulfurous  Acid,  H2S03. 

When  sulfur  dioxid  is  passed  into  water,  the  gas  is  absorbed  and  the 
solution  becomes  acid,  owing  to  an  interaction  between  the  gas  and  the 
water  to  form  a  compound  knowrn  as  sulfurous  acid.  Only  a  small  por- 
tion of  the  gas  is  actually  united  chemically  with  the  water  and  in  any 
solution  the  condition  of  equilibrium  is  one  in  which  a  large  proportion 
of  the  S02  is  simply  dissolved  in  the  solvent.  This  is  indicated  by  the 


SULFITES.  185 

fact  that,  on  boiling  or  evaporating  the  solution  spontaneously,  just  in 
proportion  as  the  gas  is  removed,  the  sulfurous  acid  is  decomposed  into 
its  anhydrid,  SO2,  and  these  reactions  continue  simultaneously  until 
finally  nothing  but  water  remains.  Inasmuch  as  sulfurous  acid  cannot 
be  isolated,  its  formula  is  not  known  with  certainty  but  it  probably  is 
H2S03.  It  acts  as  a  weak  dibasic  acid,  forming  salts  which  are  fairly 
stable.  On  the  addition  of  an  acid  to  a  sulfite,  however,  sulfur  dioxid  is 
liberated  for  the  reasons  already  set  forth. 

Na— 0  H— 0  0 

\  \  // 

S  =  0  +  2HCl-»2NaCl  +  S  =  0-*£        +H20 

Na— 0  H— 0  0 

All  sulfites  are  strong  reducing  agents,  absorbing  one  atom  of  oxygen 
for  each  molecule  of  sulfite,  forming  sulfates.  lodin,  bromin  and  per- 
manganate are  bleached  at  once  and  other  oxidizing  agents  quickly  re- 
duced. Mercurous  salts  are  unaffected  in  the  cold  but,  on  boiling,  are 
reduced  by  sulfurous  acid  exactly  as  with  stannous  chlorid.  At  ordinary 
temperatures  atmospheric  oxygen  causes  a  slow  oxidation,  especially  in 
solutions,  hence  all  sulfites  contain  some  sulfate. 

Strong  reducing  agents,  on  the  other  hand,  cause  sulfurous  acid  to 
act  as  an  oxidizing  agent.  Under  the  influence  of  nascent  hydrogen  it 
is  reduced,  first  to  hyposulfurous  acid,  H2S204,  a  weak  acid  with  strong 
reducing  properties.  The  final  product  of  the  reduction  is  H2S.  Hydro- 
gen sulfid  reduces  sulfurous  acid  and  its  anhydrid  according  to  the 
equation : 

H  0  H 

\  //  \ 

2         S  +  S  -*    2         0  +  3S 

/  \  / 

H  O  H 

Sulfites  and  other  oxidized  compounds  of  sulfur  are  reduced  to  sulfids 
by  heating  with  charcoal. 

Heated  out  of  contact  with  the  air,  sulfites  generally  decompose  into 
S02  and  metallic  oxids.  Alkaline  sulfites,  under  the  same  circumstances, 
decompose  according  to  the  equation : 

Na— 0  Na— 0          O  Na 

\  \  /  \ 

4  S  =  0    -*    3  S+S 

/  /\  / 

Na— O  ,  Na— 0          0  Na 


186  QUALITATIVE    ANALYSIS. 

Heated  in  the  air,  however,  sulfates  are  generally  formed. 

All  normal  sulfites  are  insoluble  or  soluble  with  difficulty  in  water, 
with  the  exception  of  the  salts  of  the  alkali  metals.  All  acid  sulfites  are 
soluble.  In  the  stronger  acids,  all  sulfites  dissolve  with  readiness. 

All  sulfites  are  transposed  by  boiling  with  Na2C03. 

The  barium  ion,  Ba,  precipitates  white  granular  BaS03  from  neutral 
or  alkaline  solutions  containing  the  S03  ion.  It  is  readily  soluble  in 
dilute  HC1  and  will  not  precipitate  in  the  presence  of  acetic  acid  (dis- 
tinction from  BaSOJ. 

The  silver  ion,  A+g,  precipitates  white  granular  Ag2S03,  which,  on 
boiling,  precipitates  dark  gray  metallic  silver  according  to  the  equation : 

2Ag2S03  -»  Ag2S04  +  S02  +  2Ag 

Silver  sulfite  is  soluble  in  dilute  HN03,  like  all  sulfites,  and  in 
NH4OH,  like  most  silver  salts.  It  dissolves  in  solutions  of  alkaline  sul- 
fites, forming  the  soluble  salt,  AgXaS03. 

Sulfuric  Acid,  H2S04. 

Pure  sulfuric  acid  is  a  colorless,  oily  liquid  at  ordinary  temperatures 
but,  on  cooling,  solidifies  to  a  white  crystalline  body,  melting  at  10.5°. 
When  heated,  it  decomposes  somewhat  into  its  anhydrid,  S03  and  water, 
and  it  boils  at  338°.  The  distillate  is  not  pure  H2S04,  however,  for  it 
contains  about  1.5  per  cent,  of  water  formed  during  the  heating,  for  the 
anhydrid  and  the  water  do  not  unite  completely  again  on  cooling.  The 
product  obtained  by  distillation  has  a  specific  gravity  of  1.84  and  con- 
tains about  98.5  per  cent.  H2S04.  Commercial  acid,  which  is  not  dis- 
tilled but  is  obtained  from  a  dilute  acid  by  heating  to  drive  off  as  much 
of  the  water  as  possible,  contains  only  93-96  per  cent.  H2S04.  The  prod- 
uct obtained  by  either  process  is  known  as  concentrated  sulfuric  acid. 

Fuming  sulfuric  acid  consists  of  concentrated  sulfuric  acid  in  which 
the  anhydrid,  S03,  has  been  dissolved.  From  it,  at  low  temperatures, 
pyro-sulfuric  acid,  H2S207,  can  be  crystallized  but  this  decomposes  into 
H2S04  and  SO3  on  warming. 


H— 0          0  0          H— O— S 

\//  //  I 

S         +S=O    <->  O 


/  \        \  l=o 

H— O          O  O          H— 0— S_X 


SULFUKIC    ACJD.  187 

Fuming  sulfuric  acid  has  a  greater  decomposing  action  than  concen- 
trated sulfuric  acid  because  of  the  anhydrid  it  contains. 

When  concentrated  sulfuric  acid  is  mixed  with  water,  a  decided  rise 
in  temperature  is  noted  due  to  the  formation  of  two  hydrates,  H2S04.H20 
and  H2S04.2H20.  These  hydrates  probably  are  the  normal  hexahy- 
droxid  of  sulfur  and  its  first  dehydration  product,  as  the  equations  show  : 
(See  page  174.) 


H—  O/ 

H—  0\ 

H—  0\  H—  0\ 

H~°\  O     -U     TTO  H-0\ 

= 


H— 

H—  O/  H—  O/ 

H—  O/ 

Because  of  this  reaction,  concentrated  sulfuric  acid  absorbs  water  vapor 
very  readily  and  holds  it  firmly,  hence  it  is  much  used  as  a  drying  agent. 
Sulfuric  acid  is  the  strongest  of  the  common  dibasic  acids,  dissociat- 
ing in  steps  according  to  the  rule  for  polybasic  acids  : 


(1)  H2S04-»H  +  HS04  (2)  HS04-^H+S04 

In  a  solution  of  ordinary  dilution,  the  second  step  of  the  reaction  takes 
place  to  about  the  same  extent  as  the  first  step,  so  that  the  prevalent  ion 
is  S04  and  not  HS04. 

In  dilute  solution  H2S04  acts  simply  as  an  acid  but  hot  concentrated 
sulfuric  acid  is  an  oxidizing  agent,  being  readily  reduced  to  sulfurous 
acid  which  decomposes  with  the  liberation  of  S02  and  water.  When 
heated  with  Na2C03  on  charcoal,  sulfates,  and  other  oxidized  compounds 
containing  sulfur,  are  reduced  to  sulfids.  A  similar  reduction  of  sul- 
fates to  sulfids  takes  place  in  natural  waters  under  the  influence  of  or- 
ganic matter  and  certain  bacteria  and  is  one  of  the  sources  of  "sulfur 
water,"  (H2S). 

On  ignition,  sulfates  of  the  alkali  metals  fuse  unchanged,  other  sul- 
fates are  more  or  less  decomposed  into  the  anhydrid  and  metallic  oxids 
while  the  salts  of  barium,  strontium,  calcium  and  lead  are  practically  un- 
changed. Acid  sulfates  break  down  readily,  first  into  pyro-sulfates  and, 
finally,  into  normal  sulfates  with  the  liberation  of  the  anhydrid  in  the 
nascent  condition. 


188  QUALITATIVE    ANALYSIS. 

2HNaS04  -^  H2O  +  Na2S207  ->  Na2S04  +  80 3 

Nascent  sulfur  trioxid  formed  in  this  way  is  particularly  effective  in 
decomposing  silicates  and  oxids  insoluble  in  acids.  Fusion  with  acid 
potassium  sulfate  is  often  a  satisfactory  means  of  obtaining  a  refractory 
substance  in  soluble  form.  (See  page  242.) 

Sulfates  are  generally  soluble  in  water,  the  exceptions  being  BaS04, 
SrS04,  CaS04  and  PbS04.  These  compounds,  being  salts  of  a  strong 
acid,  do  not  dissolve  to  any  great  extent  in  solutions  of  the  stronger  acids, 
HC1  or  HN03. 

Several  soluble  sulfates  separate  from  aqueous  solution  in  the  same 
crystal  form,  each  taking  up  seven  molecules  of  water  of  crystallization. 
These  compounds  are  known  as  * '  vitriols, ' '  a  term  derived  from  the  older 
name  of  sulf uric  acid, ' 1  oil  of  vitriol. ' '  The  sulfates  of  ferrous  iron,  zinc, 
magnesium,  cadmium,  cobalt,  nickel,  manganese  and  copper  are  the  best 
known  of  the  ' '  vitriols. ' ' 

All  sulfates  are  transposed  by  boiling  with  Na2C03,  with  the  excep- 
tion of  BaS04,  which  requires  fusion  with  the  same  reagent  (see  page  99) . 

The  barium  ion,  Ba,  added  to  solutions  containing  the  S04  ion,  precipi- 
tates white  crystalline  BaSO4,  practically  insoluble  in  all  solvents. 

The  silver  ion,  Ag,  gives  no  precipitate,  except  in  very  concentrated 
solutions  when  white  Ag2S04  falls. 

Thiosulfuric  Acid,  H2S208. 

Just  as  H2S03  is  oxidized  by  oxygen  to  H2S04,  so  H2S03  may  be  oxi- 
dized by  sulfur  to  H2S203,  thiosulfuric  acid. 

H— 0  H— 0  O 

\  \  // 

s=o  +  o-*          s 

/  /  \ 

H— 0  H— O          O 

H— 0  H— O          O 

\  \/ 

s=o  +  s-»          s 

/  /\ 

H— 0  H— S  O 

Although  thiosulfuric  acid  corresponds  in  its  degree  of  oxidation  to  sul- 
furic  acid,  the  presence  of  the  second  sulfur  atom  in  the  molecule  makes  it 
resemble  sulfurous  acid  far  more  than  sulfuric  acid.  While  a  whole 


THIOSULFATES.  189 

series  of  salts  are  well  known,  the  acid  itself  is  entirely  unknown,  for  it 
is  quite  as  unstable  as  sulfurous  acid  and,  when  set  free  from  one  of  its 
salts,  it  decomposes  into  S02,  water  and  free  sulfur,  the  precipitation  of 
the  latter  element  serving  to  distinguish  thiosulfates  from  sulfites.  The 
similarity  of  decomposition  products  would  naturally  suggest  a  similarity 
in  chemical  properties  and  thiosulfuric  acid  proves  to  be  a  vigorous 
reducing  agent.  Strong  oxidizing  agents,  such  as  chlorin  and  bromin, 
change  it  immediately  into  sulfuric  acid  with  or  without  a  partial  libera- 
tion of  sulfur.  Hence  the  amount  of  oxygen  taken  up  by  the  thiosulfate 
will  vary  under  different  conditions.  The  milder  oxidizing  agents,  such 
as  iodin,  do  not  carry  the  oxidation  as  far  as  H2S04,  as  the  equation 
shows : 

2H2S203  +  21  -»  H2S406  +  2HI 

Tetrathionic  acid,  H2S40G,  is  one  of  a  series  of  polythionic  acids  of 
which  the  following  members  are  known:  H2S206,  H2S306,  H2S4O6, 
H2S506.  They  are  all  strong  but  quite  unstable  acids,  readily  oxidizing 
to  sulfuric  acid.  Tetrathionic  acid  is  oxidized  by  bromin  or  other  strong 
oxidizing  agent  to  H2S04  according  to  the  equation : 

H2S406  +  14Br  +  10H20  ->  4H2S04  +  14HBr 

The  barium  salt  of  tetrathionic  acid  is  soluble,  hence  this  acid  may 
be  differentiated  from  sulfuric  acid. 

Thiosulfuric  acid  reduces  ferric  salts  to  the  ferrous  condition  and 
cupric  salts  to  the  cuprous  condition.  Cuprous  salts  are  generally  in- 
soluble, but  a  soluble  complex  sodium-cuprous  thiosulfate  is  formed  in 
this  case  by  an  excess  of  an  alkaline  thiosulfate. 

Salts  of  thiosulfuric  acid  are  much  more  stable  than  the  free  acid, 
but,  on  heating,  they  break  down  analogously  to  sulfites,  the  excess  of 
sulfur  appearing  as  a  sublimate  in  closed-tube  reactions : 

4Na2S203  ->  3Na2S04  +  Na2S  +  4S 

Heated  with  Na2C03  on  charcoal,  thiosulfates  are  reduced  to  sulfids, 
like  other  oxidized  sulfur  compounds. 

All  thiosulfates  dissolve  in  water,  only  the  barium,  lead  and  silver 
salts  being  soluble  with  difficulty. 

All  thiosulfates  are  transposed  by  boiling  with  Na2C03. 

Thiosulfuric  acid  appears  to  be  a  much  stronger  acid  than  sulfurous 
aeid.  Its  salts,  of  which  the  most  important  is  the  "hypo"  of  pho- 


190  QUALITATIVE    ANALYSIS. 

tographers,  Na2S203.5H20,  separate  the  ion,  S2O3,  with  readiness,  but,  in 
the  presence  of  H  ions,  it  is  unstable  and  decomposes  into  the  ion  S03 
and  free  sulfur.  Thiosulfates  show  a  strong  tendency  toward  the  for- 
mation of  complex  ions  with  several  metals,  as  is  shown  by  the  for- 
mation of  soluble  complex  salts,  which  do  not  give  reactions  for  the 
heavy  metal  present.  Of  these,  the  complex  sodium-silver  thiosulfate 
is  most  important  because  of  its  application  in  photography.  See  later. 

The  -barium  ion,  Ba,  precipitates  white  BaS203  from  concentrated 
solutions  only,  since  it  is  soluble  in  480  parts  of  water. 

The  silver  ion,  Ag,  if  present  in  the  correct  proportion,  will  precipi- 
tate white  Ag2S203.  This  dissolves  in  an  excess  of  an  alkaline  thiosulfate, 
forming  NaAgS203,  which  is  quite  soluble.  The  latter  salt  unites  with 
more  alkaline  thiosulfate  to  form  the  salt  Na4Ag2(S203)3,  which  is  ex- 
ceedingly soluble,  and  in  solution  gives  practically  no  Ag  ions.  In  acid 
solution,  on  boiling,  this  complex  salt  decomposes  and  silver  sulfid  is 
precipitated. 

Persulfuric  acid,  H2S208,  is  an  unstable  compound  containing  hepta- 
valent  sulfur.  Its  salts  are  much  better  known  than  the  acid.  They 
act  as  strong  oxidizing  agents,  taking  up  water  and  decomposing  readily 
into  H2S04  and  nascent  oxygen. 

Group  6,  Second  Type. 
Chromium  52.1 ;      Molybdenum  96 ;       Tungsten  184 ;      Uranium  238.5. 

The  compounds  formed  by  the  elements  belonging  to  this  second  type 
resemble  closely  those  of  the  preceding  type,  but  these  free  elements  are 
decidedly  more  metallic  than  the  former,  being  characterized  by  brilliant 
lustre,  high  specific  gravity  and  good  conductivity  of  heat  and  electricity. 
Hydrids  are  unknown  or  unimportant.  Each  element  forms  several 
oxids  and  hydroxids.  Of  the  latter,  those  in  which  the  valence  of  the 
element  is  low  (two  or  three),  possess  weakly  basic  properties,  while  those 
in  which  the  characteristic  valence  of  six,  or  a  higher  valence,  is  shown, 
are  acids  of  considerable  strength. 

Chromium  is  the  most  commonly-occurring  and,  as  a  type  of  the 
other  members  of  the  group,  will  be  considered  in  detail. 

CHROMIUM.     Cr  — 52.1. 

Chromium  is  a  difficultly  fusible,  gray-white,  crystalline  metal.  It  is 
not  changed  by  air  or  water  at  ordinary  temperatures  but,  when  heated, 
burns  to  the  green  oxid,  Cr203.  The  metal  is  very  infrequently  met  with, 
since  it  is  only  with  difficulty  reduced  from  its  compounds. 


CHROMIUM.  191 

Chromium  dissolves  fairly  well  in  dilute  HC1  and  H2S04.  Concen- 
trated acids  in  general,  and  HN03  even  when  dilute  have  but  little 
effect ;  for  the  metal  enters  the  ' '  passive ' '  condition  in  which  it  is  not  dis- 
solved by  these  reagents.  Much  energy  is  liberated  when  the  element 
passes  into  the  ionic  condition,  hence  no  other  metal  will  precipitate 
chromium  from  solutions  of  its  salts  and  it  is  only  at  very  high  tempera- 
tures that  it  can  be  reduced  in  the  * '  dry  way. ' ' 

Chromium  forms  two  well-known  series  of  compounds,  the  trivalent 
and  the  hexavalent,  and  also  two  other  series  which  are  less  familiar,  the 
divalent  and  the  heptavalent.  These  compounds  may  be  considered  as 
derived  from  four  oxids  of  which  the  first  and  fourth  are  not  known  with 
certainty. 

\ 


O 

\  // 

^=O  O  Cr=0  6 

/  \ 

Cr= O  0 


\ 


\ 
O 

Divalent  chromium  acts  as  a  fairly  strong  base-forming  element. 
Bodies  of  this  nature  may  be  formed  by  the  action  of  the  strongest  re- 
ducing agent,  i.  e.,  nascent  hydrogen,  on  compounds  in  which  the  element 
shows  a  higher  valence.  Precautions  must  be  taken  to  exclude  the 
atmosphere  entirely,  for  on  exposure  to  the  air,  they  oxidize  almost  in- 
stantly to  the  trivalent  condition.  Hence  they  are  very  seldom  encoun- 
tered and  will  not  be  further  considered. 

Trivalent  chromium  acts  as  a  weak  base-forming  element.  Its  com- 
pounds are  all  of  a  green  or  violet  color.  All  of  the  salts  hydrolyze  de- 
cidedly, several  of  the  common  ones  being  soluble.  In  neutral  or  acid 
solutions  they  are  stable,  but,  in  alkaline  solution,  they  slowly  take  up 
oxygen  and  pass  into  the  hexavalent  conditions,  the  transition  being  indi- 
cated by  a  change  in  color  from  green  to  yellow  or  orange.  Under  the 
influence  of  oxidizing  agents,  this  change  is  accomplished  very  readily, 
taking  place  more  rapidly  in  alkaline  than  in  acid  solutions.  Sodium 
peroxid,  Na202,  is  an  especially  satisfactory  oxidizing  agent  in  the  "wet 
way,"  and  KNO3  in  the  "dry  way."  A  typical  equation  might  be 
written : 


192  QUALITATIVE    ANALYSIS. 

0— H  Na— 0  0 

/  \      // 

2  Cr—  0— H  +  30  +  4NaOH    -^2  Cr        +  5H2O 

0— H  Na— 0  0 

Hexavalent  chromium  forms  compounds  of  the  type  K2Cr04  in 
which  chromium  plays  the  part  of  an  acid-forming  element.  Compounds 
of  this  class  are  highly  colored,  either  yellow  or  reddish.  They  are 
generally  very  stable,  but  in  acid  solutions  and  in  the  presence  of  reduc- 
ing agents  (alcohol  is  especially  satisfactory),  they  act  as  strong  oxidizing 
agents  (see  page  43),  parting  with  their  oxygen  and  breaking  down 
into  the  trivalent  base-forming  condition.  Strong  ignition  is  sufficient 
to  decompose  many  compounds  of  acidic  chromium  into  basic  com- 
pounds with  the  liberation  of  oxygen.  This  transition  is  marked  by  a 
change  in  color  from  yellow  or  red  to  green. 

Heptavalent  chromium  acts  as  a  strong  acid-forming  element.  Hy- 
drogen peroxid,  acting  on  acid  solutions  of  hexavalent  chromium,  gives 
rise  to  a  deep  blue  color  which  is  thought  to  be  due  to  a  heptavalent 
chromium  compound  of  the  type,  HCr04,  known  as  perchromic  acid. 
Chromium  compounds,  in  which  the  element  showTs  a  valence  of  seven, 
are  exceedingly  unstable  and  but  little  is  known  definitely  in  regard  to 
them. 

Heated  with  Na2C03  and  KN03,  all  chromium  compounds  are  oxi- 
dized to  Na2Cr04,  forming  a  yellow  melt,  readily  soluble  in  water.  After 
acidifying  slightly  with  H(C2H302),  a  drop  of  AgN03  produces  a  red 
precipitate.  This  affords  a  very  delicate  test  for  chromium. 

Heated  before  the  blowpipe  with  Na2CO3  on  charcoal,  all  chromium 
compounds  are  brought  into  the  trivalent  condition  and  leave  a  greenish 
residue,  Cr2O3. 

Heated  in  a  borax  bead,  or  sodium  meta-phosphate  bead,  all  chro- 
mium compounds  produce  a  green  color. 

The  important  ions  of  chromium  are  the  kathion,  Cr,  and  the  anions 
C?04  and  Cr207. 

The  Chromic  Kathion,  Cr. 

For  the  reactions  of  the  chromium  ion,  Cr,  solutions  of  the  nitrate 
Cr(N03)3.9H20  serve  best.  It  is  a  very  soluble  salt  and  will  melt  in  its 
own  water  of  crystallization  at  temperatures  but  little  above  the  ordinary. 

The  hydroxyl  ion,  OH,  precipitates  gelatinous,  gray-green  Cr03H3, 
a  substance  somewhat  inclined  to  form  pseudo-solutions.  On  standing, 


CHEOMIUM    SALTS.  193 

more  readily  on  heating,  it  loses  water,  forming,  finally,  green  Cr203. 
This  substance,  when  heated,  becomes  very  resistant  to  acids  and  requires 
fusion  with  acid  sulfates  or  with  Na2C03  and  KN03  to  get  it  into  soluble 
form. 

Chromium  hydroxid  acts  as  a  weak  base,  dissolving  in  acids  to  form 

+-H- 

solutions  containing  the  ion,  Cr,  or  it  may  act  as  a  weak  acid  and  dissolve 
in  strong  alkalies  to  form  solutions  containing  the  chromite  ion,  Cr03. 
In  NaOH  or  KOH,  such  action  takes  place.  Salts  of  the  ortho  type, 
Na3Cr03,  and  of  the  meta  type,  NaCr02,  are  recognized.  On  boiling  a 
solution  of  an  alkaline  chromite,  hydrolysis  and  partial  dehydration  of 
the  chromous  acid,  i.  e.,  the  chromium  hydroxid,  takes  place,  resulting  in 
the  precipitation  of  this  compound.  In  this  respect,  chromium  differs 
from  aluminum;  for  aluminum  hydroxid  is  not  precipitated  on  boiling 
a  solution  of  alkaline  aluminates  (except  ammonium  aluminate).  Fer- 
rous meta-chromite,  Fe(Cr02)2,  is  the  ordinary  ''chrome  iron  ore,"  from 
the  fusion  of  which  with  Na2C03  and  KN03,  all  salts  of  chromium  are 
derived. 

Chromium  hydroxid  is  appreciably  soluble  in  NH4OH,  forming,  not 

+-H- 

ammonium  chromite,  but  a  complex  ion,  Cr  (NH3)4. 

Like  aluminum  and  ferric  hydroxids,  in  the  presence  of  organic  mat- 
ter chromium  hydroxid  is  partially  held  up  and  prevented  from  pre- 
cipitating. 

The  carbonate  ion,  C03,  from  solutions  of  alkaline  carbonates,  pre- 
cipitates Cr(OH)3._ 

The  sulfid  ion,  S,  in  acid  solutions  does  not  cause  a  precipitate,  for 
Cr2S3,  formed  only  in  the  "dry  way,"  is  soluble  in  acids  and  is  hydrolyzed 
completely  by  water.  By  solutions  of  alkaline  sulfids,  Cr(OH)3  is 
precipitated. 

The  sulfate  ion,  S04,  causes  no  precipitate,  for  Cr2(S04)3.15H20  is 
very  soluble.  Like  aluminum,  it  forms  double  salts,  known  as  "chrome 
alums, ' '  with  the  sulfates  of  univalent  metals.  Potassium  chrome  alum, 
Cr2(SOJ3.K2S04.24H2O,  or  better  KCr(S04)2.12H2O,  is  a  stable  violet- 
colored  salt,  soluble  in  6  or  7  parts  of  water.  Any  complex  ions  which 
may  exist  in  highly  concentrated  solutions  are  very  unstable  and,  in  dilute 
solutions,  only  the  Cr  and  other  simple  K  and  S04  ions  are  present. 

Two  modifications  of  many  soluble  chromium  salts  are  recognized.  A 
series  of  violet-colored  salts  exist,  from  the  solutions  of  which  all  the 
metal  and  all  the  acid  can  be  precipitated  by  the  ordinary  reagents,  thus 
showing  that  the  regular  ions  are  present.  By  heating  the  violet  salt,  a 
green  modification  is  formed,  from  solutions  of  which  only  a  fraction  of 
*  14 


194  QUALITATIVE    ANALYSIS. 

the  metal  and  acid  can  be  precipitated  by  the  ordinary  reagents.  This 
indicates  the  formation  of  complex  ions  from  the  simple  ions,  apparently 
by  some  sort  of  ' '  doubling  up ' '  process. 

The  chlorid  ion,  Cl,  produces  no  precipitate,  for  green  CrCl3.6H20  is 
very  soluble,  as  are  all  the  other  halids  of  chromium.  It  is  an  interest- 
ing fact  that  on  driving  off  the  water  of  crystallization  from  this  green 
salt,  it  turns  violet  as  it  dehydrates  and  becomes  almost  absolutely  insolu- 
ble in  water. 

The  cyanid  ion,  CN,  precipitates  Cr(OH)3. 

The  phosphate  ion,  P04,  precipitates  gray-green  CrP04,  soluble  in 
dilute  mineral  acids  but  not  in  acetic  acid.  In  this  respect,  it  resembles 
aluminum"  and  ferric  phosphates. 

The  acetate  ion,  (C2H302),  produces  no  precipitate,  as  Cr(C2H3C2)3 
is  soluble.  Whether  or  not  chromium  forms  a  basic  acetate,  like  alu- 
minum and  ferric  iron,  is  undecided.  If  a  solution  of  chromium  acetate 
is  boiled  by  itself,  no  precipitate  forms.  If  considerable  ferric  acetate 
be  present,  the  chromium  is  found  practically  completely  precipitated  in 
some  form  along  with  the  basic  ferric  acetate.  It  is  necessary,  therefore, 
to  insure  the  presence  of  ferric  salts  before  attempting  to  precipitate 
chromium  in  this  way.  (See  page  264.) 

The  Chromate  Ion,  Cr04,  and  the  Bichromate  Ion,  Cr207. 

Chromium  trioxid,  Cr03,  is  a  red  crystalline  substance,  readily  solu- 
ble in  water,  forming  orange-red  solutions  which  have  a  decidedly  acid 
reaction.  On  making  neutral  or  alkaline,  the  color  changes  to  yellow  and 
there  crystallizes  out  yellow  salts  of  the  type,  K2Cr04,  which  are  known 
as  chromates.  Such  yellow  solutions  contain  the  chromate  ion,  Cr04. 
which  is  similar  to  the  sulfate  ion,  S04.  On  acidifying  a  solution  of  a 
chromate,  the  color  becomes  orange-red  and  a  red  dichromate  crystallizes, 
of  which  the  potassium  salt,  K2Cr2O7,  may  be  taken  as  typical.  The 
orange-red  solutions  contain  the  dichromate  ion,  Cr207. 

The  chromate  ion,  CrO4,  is  unstable  in  the  presence  of  H  ions  while 
the  Cr207  ion  is  unstable  in  the  presence  of  OH  ions.  The  change  from 
the  one  into  the  other  is  expressed  by  the  equations : 

2CrO4  +  2H  ->  CrA  +  H20 
O2O7  +  2  OH  -»  2Cr04  +  H20 

For  this  reason,  any  attempt  to  obtain  chromic  acid,  H2Cr04,  is 
unavailing  and  results  only  in  dichromic  acid,  H2Cr2O7.  Even  this  acid 


CHEOMIC    AND    DICHROMIC    ACIDS.  195 

is  known  only  in  solution  and  every  attempt  to  isolate  it  results  in  the 
anhydrid,  CrO3,  according  to  the  equation : 

H2O207  ->  2O03  +  H2O 

Dichromic  acid  may  be  derived  theoretically  by  separating  one  mole- 
cule of  H2O  from  two  molecules  of  H2Cr04 : 

0 

H— 0  0  H— 0— Cr=0 

\     /  \ 

2  Cr  ->  0     +     H2O    • 

/     \  / 

H— 0  0  H— 0— Cr— 0 

\ 
O 

Dichromic  acid  is,  therefore,  entirely  analogous  to  other  acids  of  the 
pyro-  type  and  is  to  be  considered  as  pyro-chromic  acid.  Since  the  dif- 
ference between  the  acids  is  merely  a  question  of  dehydration,  it  would 
seem  as  though  dichromic  acid  in  dilute  solution  might  become  hydrated 
and  changed  into  chromic  acid,  in  a  measure  at  least,  by  the  mass  action 
of  the  water.  The  same  considerations  would  seem  to  indicate  that' a 
dichromate  in  dilute  solutions  might  be  changed  into  a  chromate,  partially 
at  least.  As  a  matter  of  fact  this  is  so ;  for,  if  into  solutions  containing 
dichromate  ions  various  kathions  be  introduced,  metallic  dichromates 
are  seldom  formed.  In  all  dilute  solutions  chromates  alone  are  formed, 
whether  the  original  solution  contained  Cr707  or  Cr04  ions.  This  phe- 
nomenon is  largely  explained  by  the  fact  that  few  dichromates  are  known 
except  salts  of  the  alkali  and  alkaline  earth  metals.  Where  both  types  are 
formed  chromates  are  much  less  soluble,  as  a  rule,  than  the  corresponding 
dichromates  and,  therefore,  are  precipitated.  The  following  equations 
illustrate  this  reaction : 

2Pb  ( C2H302)  2  +  K2Cr207  +  H20  -» 

2PbCr04  +  2K(C2H302)  +  2H(C2H302) 

or  writing  an  ionic  equation : 

2Pb  +  Cr7O7  +  H2O  -*  2PbCrO4  +  2H 
Chromates  show  a  close  similarity  to  sulf  ates  in  many  of  their  proper- 


196  QUALITATIVE    ANALYSIS. 

ties,  and  a  comparison  would  seem  to  indicate  that  the  hypothetical 
chromic  acid  is  somewhat  weaker  than  sulfuric  acid.  Chromates  are 
readily  soluble  in  the  stronger  acids,  therefore,  but  do  not  dissolve  in 
acetic  acid.  In  high  concentration,  however,  this  last  reagent  is  able 
to  prevent  the  formation  of  a  precipitate,  although  it  is  usually  not 
strong  enough  to  dissolve  the  salt  when  once  brought  down. 

The  alkali  chromates  are  all  very  soluble,  K2O04  being  the  most  com- 
mon salt.  It  dissolves  in  less  than  2  parts  of  H20  at  ordinary  tempera- 
tures and  separates  from  solution  without  water  of  crystallization.  The 
ammonium  salt  is  similar  to  the  potassium  salt.  The  sodium  salt, 
Na2Cr04.10H20,  is  very  hydroscopic  and  melts  in  its  own  water  of  crys- 
tallization at  temperatures  but  little  above  the  ordinary.  Because  of  the 
great  solubility  of  these  salts,  they  are  ordinarily  used  as  the  reagents  for 
introducing  Cr04  ions  into  solutions  of  other  salts. 

Most  chromates  are  soluble  in  water,  the  exceptions  being  the  salts 
of  silver,  mercury  in  the  lower  condition,  lead,  copper,  bismuth,  tin  and 
antimony  in  the  lower  condition,  cobalt,  nickel  and  barium. 

All  chromates  are  transposed  by  boiling  with  Na2C03. 

Of  the  dichromates,  salts  of  the  alkalies  and  alkaline  earths,  together 
with  silver,  mercury,  lead  and  copper  only,  are  well  known.  These  are 
soluble  or  are  decomposed  by  water. 

Heated  with  concentrated  H2S04,  chromates  or  dichromates,  mixed 
with  chlorids,  give  rise  to  a  dark,  reddish  liquid,  boiling  at  117° 
and  known  as  chloro-chromic  anhydrid.  Since  the  corresponding  bromin 
and  iodin  compounds  are  not  formed  under  the  same  conditions,  this 
offers  a  means  of  differentiating  between  salts  of  the  halogen  elements 
(see  page  299).  The  reaction  by  which  chloro-chromic  anhydrid  is 
formed  may  be  represented  by  the  equation : 

0 

//Cl 

K2Cr04  +  2KC1  +  2H2S04  -*  Cr<      +  2K2S04  +  2H20 

\Cl 

0 

This  equation  represents,  however,  only  the  initial  and  final  stages  in 
the  reaction  and  the  true  function  of  the  concentrated  H2S04  is  lost 
sight  of.  A  better  idea  of  the  process  may  be  obtained  by  a  consideration 
of  the  following  equations,  which  indicate  the  steps  in  the  reaction : 


CHEOMATES.  197 

K2Cr04  +  H2S04  -*  K2S04  +  H2Cr04     2KC1  +  H2S04  ->  K2S04  +  2HC1 

H2Cr04  -»  H20  +  O03 
O  0 

//         HC1  //Cl 

Cr=0  +  -»    Cr<      +H2O 

\         HC1  \CZ 

O  0 

As  will  be  noted,  water  is  liberated  during  the  process  of  formation 
of  Cr02Cl2,  but  is  prevented  from  decomposing  the  latter  by  the  large 
excess  of  H2S04  which  is  present.  If,  however,  water  is  present  in  excess, 
it  reacts  with  the  chloro-chromic  anhydrid  exactly  as  it  does  with  halogen 
compounds  of  non-metallie,  i.  e.,  acid-forming  elements  generally.  It  is 
hydrolyzed  with  the  formation  of  two  acids,  for  hexavalent  chromium 
is  an  acid-forming  element. 

0  0 

/Cl      HOH  /OH 

Cr<      +  -»     Cr<        +2HC1 
\C1      HOH  \OH 

0  0 

The  barium  ion,  Ba,  added  to  dilute  solutions  containing  chromates  or 
dichromates,  precipitates  canary-colored  barium  chromate,  BaCr04,  in- 
soluble in  all  alkalies.  It  is  soluble  in  about  250,000  parts  of  water. 

The  silver  ion,  Ag,  precipitates  red  silver  chromate,  Ag2Cr04,  readily 
soluble  in  NH4OH  (see  page  86).  It  dissolves  in  about  40,000  parts  of 
H20. 

The  lead  ion,  Pb,  precipitates  deep  yellow  lead  chromate,  PbCr04, 
soluble  in  the  strong  alkalies  (see  page  145).  It  is  very  highly  insoluble 
in  pure  water,  requiring  several  hundred  thousand  parts  of  water  for  its 
solution.  It  is,  however,  much  more  soluble  in  traces  of  acids,  bases  and 
other  salts. 

The  mercurous  ion,  (Hg)2,  precipitates  a  brown  basic  chromate  which, 
on  foiling,  changes  to  the  fiery-red  normal  chromate,  Hg2Cr04,  insoluble 
in  alkalies, 

Group  7,  First  Type. 
Fluorin  19 ;         Chlorin  35.5 ;         Bromin  80 ;         lodin  127. 

The  members  of  this  most  non-metallic  and  chemically-active  group 
are  known  as  halogens  or  "salt  formers,"  sea-salt,  NaCl,  being  a  very 
common  substance  derived  from  one  of  the  members  of  this  group.  The 


198  QUALITATIVE    ANALYSIS. 

most  characteristic  valence  is  one ;  with  the  exception  of  fluorin,  valences 
of  five  or  seven,  sometimes  three,  are  shown  by  these  elements  in  well- 
known  compounds,  however.  The  hydrids,  of  the  type  HC1,  possess  acid 
properties  in  a  marked  degree,  hydrofluoric  acid  being  considerably 
weaker  than  the  others.  The  oxids  of  these  elements  are  acid  anhydrids 
but,  with  the  exception  of  the  iodin  oxids,  they  generally  possess  little 
stability,  being  often  explosive  or  entirely  unknown.  The  acids  formed 
by  the  union  of  these  anhydrids  with  water  are  generally  inclined  to 
ready  decomposition  also.  The  stability  as  well  as  the  strength  of  the 
acid  increases  as  the  valence  of  the  halogen  rises,  however,  and  hepta- 
valent  compounds  of  the  perchloric  acid  type,  HC104,  are  very  strong  as 
well  as  very  stable. 

FLUORIN.    F  — 19. 

Fluorin  is  a  pale  yellowish-green  gas  with  a  very  pungent  odor 
and  toxic  properties.  It  is  the  most  reactive  of  all  known  substances. 
With  the  exception  of  oxygen  and  the  inert  argon  group,  it  forms 
compounds  with  all  the  elements,  many  of  which  take  fire  spon- 
taneously in  it.  Even  at  ordinary  temperatures,  in  the  presence  of  mois- 
ture, it  unites  with  practically  every  substance,  hence  it  does  not  exist 
free  in  nature,  nor  can  it  be  preserved  for  a  long  time.  For  this  reason, 
also,  its  isolation  from  its  compounds  was  delayed  for  a  long  time. 
Water  is  decomposed  with  explosive  violence  even  in  the  dark,  hydrogen 
fluorid  being  formed,  and  oxygen,  which  is  very  rich  in  ozone,  being  set 
free.  Chlorids,  bromids  and  iodids  are  decomposed  with  the  production 
of  fluorids  and  the  liberation  of  the  halogen.  (See  similar  action  of 
chlorin,  page  200.)  Gold  and  platinum  are  not  attacked  by  pure  dry 
fluorin. 

Fluorids  may  be  considered  as  derived  from  hydrofluoric  acid  and 
their  properties,  as  well  as  the  reactions  of  the  fluorid  ion,  F,  will  be 
treated  under  this  head. 

Hydrofluoric  Acid,  HF. 

Hydrogen  fluorid  is  a  colorless  liquid  (M.  P.— 92,  B.  P.  19.5°)  with 
a  strong  pungent  odor  and  very  poisonous  properties.  It  is  very  soluble 
in  water,  more  than  four  hundred  volumes  of  gas  dissolving  in  one  vol- 
ume of  water.  A  solution  containing  35.5  per  cent,  of  HF  boils  at  120° 
and  is  known  as  * '  concentrated  hydrofluoric  acid. ' ' 

Hydrofluoric  acid  is  a  very  corrosive  liquid.  It  attacks  all  metals, 
except  gold  and  platinum  (lead  only  slightly)  ;  all  oxids ;  many  salts,  and 
other  substances  of  a  different  nature.  Since  glass  is  readily  dissolved 
by  hydrofluoric  acid,  the  latter  must  be  preserved  in  containers  made  of 
paraffine,  rubber,  or  a  few  other  materials  which  are  not  attacked  by  it. 


> 


FLUORIDS.  199 


causes  very  painful,  poisonous,  slow-healing  burns  when  in  contact 
with  the  flesh. 

The  etching  of  glass  is  one  of  the  most  important  reactions  of  hydro- 
fluoric acid.  Glass  is  essentially  a  mixture  of  several  salts  of  the  various 
polysilicic  acids.  The  composition  of  soda  glass  may  be  represented  by 
the  formula,  Na2Si307.CaSi307,  and  the  action  of  hydrofluoric  acid  by  the 
equation  : 


Na2Si307.CaSi307  +  28HF  ->  2NaF  +  CaF2  +  eSiF^  +  14H20 

A  very  delicate  test  for  hydrofluoric  acid  depends  upon  this  reaction, 
(See  page  278.) 

Hydrofluoric  acid  is  usually  considered  as  a  monobasic  acid  and,  as 
such,  ought  not  to  form  true  acid  salts  (see  page  21).  Salts  of  this 
character  are  well-known,  however.  An  investigation  of  the  molecular 
weight  of  hydrogen  fluorid  indicates  that,  at  100°  and  above,  the  molecule 
is  HF  but  that,  at  ordinary  temperatures,  it  doubles  up  and  has  the  for- 
mula (HF)2.  Acid  salts  such  as  HNaF2  and  HKF2  are  readily  ex- 
plained by  some  such  formulae  as  the  following  : 

H—  F  H—  F 

I         or  || 

Na—  F  Na—  F 

Hydrogen  fluorid  acts  as  a  comparatively  weak  acid,  dissociating  to 
an  extent  much  less  than  the  corresponding  .  acids  formed  by  the  other 
halogen  elements.  In  solution,  fluorids  are  hydrolyzed  to  some  extent 
and  glass  containers  are  etched  by  the  liberated  hydrofluoric  acid  on  long 
standing. 

On  ignition,  fluorids  fuse  unchanged,  some  of  them  being  volatile  at 
temperatures  easily  attainable. 

Heated  with  concentrated  H2S04,  all  fluorids  are  changed  to  sulfates 
and  hydrofluoric  acid  is  volatilized. 

Salts  of  hydrofluoric  acid  differ  in  their  solubility  reactions  very 
much  from  the  corresponding  chlorids,  bromids  and  iodids.  Most  fluor- 
ids are  soluble  ;  the  copper,  lead,  zinc  and  ferric  salts  dissolve  sparingly 
and  the  salts  of  the  alkaline  earths  are  quite  insoluble. 

All  fluorids  formed  by  precipitation  are  readily  transposed  by  boiling 
with  Na2C03. 

The  simple  fluorid  ion  is  F.  Several  complex  Jons  containing  fluorin 
are  recognized.  Of  these  the  silicofluorid  ion,  SiF6,  is  tlae  best  known. 
(See  page  136.) 


200  QUALITATIVE    ANALYSIS. 

++ 
The  barium  ion,  Ba,  added  to  a  solution  containing  the  fhiorid  ion, 

causes  a  voluminous  white  precipitate,  BaF2,  soluble  in  large  amounts 
of  the  stronger  acids  and  in  ammonium  chlorid. 

The  calcium  ion,  Ca,  precipitates  white  CaF2,  which  is  somewhat  less 
soluble  than  the  corresponding  barium  salt.  The  naturally-occurring 
compound,  fluorspar,  is  however,  one  of  the  most  insoluble  substances 
known.  In  order  to  get  it  completely  into  solution,  it  is  usually  neces- 
sary to  fuse  it  with  Na2C03,  to  which  some  silica  or  a  silicate  has  been 
added.  Silica  aids  in  the  decomposition  of  the  calcium  fluorid  by  the 
formation  of  calcium  silicofluorid  as  an  intermediate  product  which  is 
later  decomposed  by  the  carbonate. 

The  silver  ion,  Ag,  causes  no  precipitation,  for  silver  fluorid  is  readily 
soluble. 

CHLORIN.    Cl  — 35.5. 

Chlorin  is  a  yellowish-green  gas  with  a  suffocating  and  nauseous  odor. 
It  is  somewhat  soluble  in  water,  two  volumes  of  gas  forming  a  solution 
which  is  known  as  "chlorin  water,"  from  which  all  chlorin  may  be  re- 
moved by  boiling.  At  zero,  chlorin  unites  with  water  to  form  a  crystal- 
line hydrate  of  the  composition  C12.8H20,  which  at  higher  temperatures 
breaks  down  again  into  its  components.  The  element  is  also  soluble  in 
carbon  disulfid,  chloroform  and  ether,  giving  colorless  solutions. 

Chlorin  is  a  typical  non-metallic  element  and  is  surpassed  in  chemical 
activity  by  fluorin  only.  It  combines  readily  with  all  metals  and  with 
many  of  the  non-metallic  elements,  often  causing  them  to  burst  into  spon- 
taneous combustion.  It  unites  with  hydrogen,  to  form  hydrogen  chlorid, 
so  vigorously  that  it  will  abstract  this  element  from  many  of  its  com- 
pounds. Even  the  very  stable  body,  water,  is  slowly  decomposed  with 
the  liberation  of  oxygen.  Bromids  and  iodids  in  solution  are  changed  to 
chlorids  and  the  halogen  is  set  free. 

Chlorin  passes  with  the  greatest  readiness  into  the  ionic  condition, 
forming  the  simple  ion,  Cl.  On  the  basis  of  this  property,  the  decompo- 
sition of  water  and  of  bromids  and  iodids  in  solution  is  to  be  explained, 
for  chlorin  shows  a  greater  tendency  to  assume  the  electric  charge  than 
oxygen,  bromin  or  iodin  do  to  hold  it  (see  page  327).  Hence,  the  charge 
passes  from  these  elements,  which  are  thus  set  free,  and  is  taken  up  by 
the  chlorin,  which  becomes  the  chlorin  ion,  as  the  equations  show : 

0  or      2  OH  +  2Cl-»  2C1  +  H20  +  0 

1  +  Br  T+C1-»C1 


CHLOKIN.  201 

Chlorin  is  one  of  the  strongest  oxidizing  agents,  carrying  substances 
generally  to  the  highest  degree  of  oxidation  capable  under  the  conditions 
of  the  experiment.  Its  action  on  water  establishes  its  oxygen  equivalent, 
two  atoms  of  chlorin  being  equivalent  to  one  atom  of  oxygen,  as  the  equa- 
tion shows  : 

H0-»2HCl      0 


Chlorin  possesses  very  strong  bleaching  properties,  litmus,  turmeric, 
indigo  and  other  organic  dye-stuffs  being  immediately  destroyed  by  the 
moist  gas  or  its  solution  in  water.  The  dry  gas  is  without  appreciable 
effect,  however,  and  the  bleaching  action  of  moist  chlorin  is  apparently 
to  be  attributed  to  nascent  oxygen  set  free  by  the  decomposition  of  the 
water  present. 

Chlorin  acts  on  cold  dilute  solutions  of  the  strong  bases  to  form 
chlorids  and  hypochlorites  in  equal  proportions,  as  the  equation  indi- 
cates : 

2C1  +  2KOH  ->  KC1  +  KOC1  +  H20 

If  the  solution  be  heated  or  is  concentrated,  chlorates  (along  with 
chlorids)  are  formed  instead  of  hypochlorites,  owing  to  the  decomposi- 
tion of  the  latter  (see  page  204).  If  alkaline  carbonates  are  substituted 
for  the  hydroxids  in  the  above  reactions,  the  same  products  are  formed 
and  carbon  dioxid  is  driven  out. 

With  oxygen,  chlorin  does  not  unite  directly  but,  by  the  decomposi- 
tion of  certain  oxygen  acids  of  chlorin,  three  oxids  can  be  obtained  hav- 
ing the  composition  indicated  by  the  formulae  —  C120,  C102  and  C1207. 
These  bodies  will  be  considered  under  the  respective  acids  from  which 
they  are  derived. 

United  with  oxygen,  chlorin  forms  several  anions,  of  which  the  most 
important,  the  hypochlorite  ion,  OC1,  the  chlorate  ion,  C103,  and  the 
perchlorate  ion,  C104,  will  be  treated  under  the  respective  acids  from 
which  they  are  derived.  Inasmuch  as  the  compounds  formed  by  the 
simple  chlorin  ion,  Cl,  may  be  considered  as  salts  of  hydrochloric  acid, 
the  reactions  of  this  ion  and  the  properties  of  the  salts  will  be  considered 
under  this  head. 

Free  chlorin  is  usually  detected  by  its  odor,  bleaching  properties  and 
action  on  bromids  and  iodids.  These  reactions  serve  to  distinguish  it 
from  all  reagents  except  hypochlorous  acid,  which  has  bleaching  proper- 
ties, the  same  action  on  bromids  and  iodids  and  an  odor  closely  resem- 
bling chlorin.  Metallic  mercury  offers  a  satisfactory  means  of  differen- 
tiation, for  chlorin  forms  white  mercurous  chlorid,  (HgCl)2,  insoluble  in 


202  QUALITATIVE    ANALYSIS. 

HC1,  whereas  hypochlorous  acid  forms  brown  mercuric  oxychlorid, 
Hg2Cl20,  which  is  readily  soluble  in  HC1.  Hydrochloric  acid  is  without 
effect  on  mercury ;  therefore,  the  solution  to  be  tested  is  slightly  acidified 
with  HC1  and  shaken  with  mercury  until  it  no  longer  turns  a  drop  of 
starch-iodid  solution  blue.  A  brown  precipitate  indicates  hypochlorous 
acid,  a  white  precipitate  free  chlorin.  In  case  the  indication  is  doubtful, 
the  solution  may  be  decanted  and  the  precipitate  treated  with  HC1  when 
the  brown  oxychlorid  will  dissolve,  leaving  behind  white  mercurous 
chlorid  if  free  chlorin  was  present.  Hydrogen  sulfid  passed  into  the  HC1 
extract  will  produce  a  black  precipitate  of  HgS  if  hypochlorous  acid  was 
present. 

Hydrochloric  Acid,  HC1. 

Hydrogen  chlorid  is  a  colorless  gas  with  a  sour  taste  and  sharp  pun- 
gent odor.  It  fumes  in  the  air  and  dissolves  abundantly  in  water,  450 
volumes  of  gas  requiring  only  one  volume  of  water  for  solution.  Its 
aqueous  solution  reacts  strongly  acid  to  litmus  and  is  known  as  hydro- 
chloric acid.  ''Concentrated  hydrochloric  acid"  is  a  solution  which  is 
about  saturated  at  ordinary  temperatures  and  contains  38—39  per  cent. 
HC1  and  61-62  per  cent.  H20.  A  solution  containing  20  per  cent.  HC1 
boils  constant  at  110°,  hence  hydrogen  chlorid  cannot  be  removed  com- 
pletely from  solution  by  boiling ;  for,  on  heating,  any  solution  loses  mainly 
either  water  or  hydrogen  chlorid  until  the  composition  of  the  liquid  in  the 
vessel  becomes  20  per  cent.  HC1  and  80  per  cent.  H20.  This  solution  then 
distills  over  without  change.  (See  similar  action  of  HN03,  page  150.) 

Hydrogen  chlorid  is  a  very  stable  compound  and  in  dilute  solutions 
does  not  yield  readily  to  oxidizing  or  reducing  agents.  Strong  oxidizing 
agents  such  as  HMn04,  HC103,  HC10,  HN03,  HN02,  H2Cr207  and  H202, 
however,  oxidize  it  to  free  chlorin  and  water,  especially  in  concentrated 
solutions,  although  in  dilute  solution  the  action  is  often  inappreciable. 
When  a  dry  chlorid  is  mixed  with  powdered  K2O207  and  heated  with 
concentrated  H2S04,  the  chlorin  first  liberated  by  the  reaction  unites  with 
chromium  trioxid,  another  reaction  product,  to  form  Cr02Cl2,  a  reddish- 
brown  gas  under  the  conditions  of  this  experiment.  This  reaction  is 
often  used  as  a  test  for  HC1.  (See  pages  196  and  299.) 

Hydrochloric  acid  acts  as  one  of  the  very  strongest  monobasic  acids, 
in  solutions  of  ordinary  dilution  being  practically  completely  ionized. 

Salts  of  hydrochloric  acid  are  soluble,  as  a  rule.  The  chlorids  of 
silver  and  of  mercury,  copper,  gold  and  platinum  in  the  lower  form  are 
highly  insoluble,  however.  Lead  chlorid  may  be  precipitated  from  cold 
solutions  containing  above  one  per  cent,  of  lead  but  it  is  readily  soluble 


CHLOEIDS.  203 

in  hot  water.  The  commonly-occurring  oxychlorids  of  bismuth  and  anti- 
mony are  also  insoluble. 

On  ignition  in  the  air,  most  metallic  chlorids  fuse  unchanged  and 
volatilize  at  a  higher  temperature.  Mercuric  chlorid  melts  at  265°  and 
boils  at  307°,  lower  than  other  metallic  chlorids.  When  a  metal  forms 
more  than  one  chlorid,  the  higher  form  generally  breaks  down  into  the 
lower.  The  chlorids  of  gold  and  platinum  lose  all  their  halogen  and 
decompose  into  the  metals.  Chlorids  of  trivalent  metals  are  generally 
changed  to  the  corresponding  oxids  with  the  liberation  of  the  halogen. 

Heated  with  concentrated  H2S04,  all  chlorids  are  changed  to  sulf  ates 
and  HC1  is  liberated.  This  reaction  is  not  to  be  explained  by  the  as- 
sumption that  H2S04  is  stronger  than  HC1  (for  the  converse  is  true),  but 
it  is  entirely  due  to  the  fact  that  HC1  is  much  more  volatile  than  H2S04 
and  is  therefore  driven  off.  (See  page  27.) 

All  chlorids  are  transposed  by  boiling  with  Na2C03,  except  silver 
chlorid,  which  requires  fusion  with  the  dry  reagent.  Silver  carbonate, 
the  first  reaction  product,  breaks  down  into  the  metal,  carbon  dioxid  and 
oxygen. 

The  barium  ion,  Ba,  added  to  a  solution  containing  the  chlorid  ion, 
01,  causes  no  precipitate,  for  BaCl2  is  readily  soluble. 

The  silver  ion,  Ag,  precipitates  curdy  white  AgCl,  darkening  in  the 
sun-light.  (See  page  86.) 

-H- 

The  lead  ion,  Pb,  precipitates  white  PbCl2  from  cold  solutions  which 
are  not  too  dilute.  (See  page  145.) 

The  mercurous  ion,  (Hg)2,  precipitates  white  (HgCl)2.  (See  page 
111.) 

Hypochlorous  Acid,  HC10. 

Salts  of  hypochlorous  acid  may  be  produced  by  the  action  of  chlorin 
on  alkalies  or  alkaline  carbonates,  as  has  been  previously  stated.  If 
nitric  or  sulfuric  acid  be  added  to  such  a  solution,  hypochlorous  acid  is 
liberated  but,  on  every  attempt  to  isolate  it,  its  decomposition  products 
only,  chiefly  chlorin,  oxygen  and  hydrochloric  acid,  are  obtained.  This 
is  due,  partially  to  the  formation  of  the  anhydrid,  C120,  an  exceedingly 
unstable  substance  which  decomposes  immediately  into  its  elements,  and 
partially  to  a  general  reaction  of  all  hypochlorites.  When  in  concen- 
trated solution  or  when  warmed,  they  undergo  a  reaction,  forming  chlor- 
ids and  chlorates  in  equal  proportions,  as  the  equation  shows : 


204  QUALITATIVE    ANALYSIS. 

0 

// 

3K—  0—  Cl    ->    2K—  Cl     +     K—  0—  Cl 

\ 
0 

Hypochlorous  acid  acts  as  a  monobasic  acid  but  it  is  so  weak  that  the 
carbon  dioxid  of  the  atmosphere  liberates  it  from  its  salts.  Hence  it  is 
that  hypochlorites,  exposed  to  the  air,  are  soon  decomposed  entirely. 
Further  indication  of  the  weakness  of  hypochlorous  acid  is  furnished  by 
the  fact  that  solutions  of  the  salts  formed  by  the  alkali  metals  react 
strongly  basic  because  of  hydrolysis. 

All  hypochlorites  possess  strong  bleaching  properties.  This  is  ex- 
plained by  the  fact  that  they  are  vigorous  oxidizing  agents.  In  the  pres- 
ence of  reducing  matter,  each  molecule  of  hypochlorite  gives  up  one  atom 
of  oxygen  and  becomes  a  chlorid.  Bleaching  takes  place  in  alkaline  as 
well  as  in  neutral  and  acid  solutions,  a  reaction  serving  to  differentiate 
hypochlorites  from  other  compounds  of  chlorin.  The  bleaching  action  is 
most  rapid  in  acid  solution  because  of  the  greater  instability  of  the  free 
acid  than  of  the  salts.  i  '  Chloride  of  lime,  '  '  the  great  commercial  bleach- 
ing salt,  is  a  mixture  of  calcium  chlorid  with  the  hypochlorite,  made  by 
passing  chlorin  over  slacked  lime. 

All  hypochlorites  are  accompanied  by  chlorids  because  of  the  method 
of  their  formation,  hence  when  any  ordinary  acid  is  added  to  a  solution 
of  a  hypochlorite,  hydrochloric  acid,  as  well  as  hypochlorous  acid,  is 
liberated.  These  reagents  interact  ;  the  hydrochloric  acid  is  oxidized  and 
chlorin  becomes  the  chief  decomposition  product  under  these  conditions. 

HC1  +  HOC1  -*  H20  +  C12 

Inasmuch  as  chlorin  water  cannot  be  preserved  for  any  considerable 
time,  an  acidified  solution  of  a  hypochlorite  is  often  used  instead  of  the 
other  more  dilute  reagent. 

All  hypochlorites  are  soluble  and  may  be  transposed  by  boiling  with 
Na2C03.  The  ammonium  salt  is  unknown,  for,  on  bringing  together 
ammonia  and  hypochlorous  acid,  as  fast  as  the  ammonium  hypochlorite 
is  formed,  it  interacts  with  the  remaining  ammonia  and  is  decomposed, 
according  to  the  equation: 

3NH4C10  +  2NH3  -*  3NH4C1  +  3H20 


+4- 

The  barium  ion,  Ba,  added  to  a  solution  containing  the  hypochlorite 
ion,  CIO,  causes  no  precipitation. 


HYPOCHLOKITES.  205 

The  silver  ion,  A+g,  precipitates  white  AgCl,  owing  to  the  almost  im- 
mediate decomposition  of  the  soluble  silver  hypochlorite  according  to  the 
regular  equation : 

3AgC10  -»  AgC103  +  2AgCl 

One-third  of  the  silver  remains  in  solution  as  the  chlorate,  however, 
as  the  equation  shows. 

Chloric  Acid,  HC103. 

Although  chloric  acid  is  considerably  more  stable  than  hypochlorous 
acid,  like  the  latter  it  is  known  only  in  solution  and  attempts  to  isolate 
it  give  only  decomposition  products: 

O  00 

/  //          // 

3H— 0— Cl        ->H— 0—01=0     +     2C1        +  H20 

\  \  \ 

0  00 

On  adding  an  acid  to  a  chlorate,  therefore,  C102  is  an  important  de- 
composition product. 

Chlorin  dioxid  is  a  heavy  yellowish-brown  gas  with  stifling  odor.  It 
often  decomposes  spontaneously  with  explosive  violence  into  the  ele- 
ments. Its  solution  in  water  is  more  stable  and,  after  a  time,  takes  an 
acid  reaction.  After  neutralizing  with  an  alkali,  evaporation  of  the 
solution  leaves  a  residue  consisting  of  a  chlorate  and  a  chlorite.  The 
corresponding  chlorous  acid,  HC102,  is  unknown  except  in  dilute  solu- 
tion, and  its  salts  are  of  little  importance.  (See  action  of  N02,  page  148.) 

O  0 

/  / 

2C1         +     H20    -»    H— 0— 01=0     +     H— 0— Cl 

\  \ 

0  0 

Chlorin  dioxid,  being  so  unstable  and  breaking  down  into  nascent 
oxygen  and  chlorin,  is  one  of  the  most  vigorous  oxidizing  agents  known. 
For  this  reason,  a  solution  of  chloric  acid  has  strong  oxidizing  and  bleach- 
ing properties  but,  unlike  hypochlorites,  chlorates  bleach  only  in  acid 
and  not  in  neutral  or  alkaline  solutions;  for  chlorates,  in  solution,  are 
stable  and  only  chloric  acid  decomposes  in  the  manner  stated. 

When  a  chlorate  is  acidified  with  HC1,  much  of  the  latter  is  oxidized 
by  the  free  chloric  acid,  which  is  liberated,  or  by  its  decomposition  prod- 
uct, C102.  Free  chlorin  is  liberated  in  abundance,  therefore. 


206  .  QUALITATIVE    ANALYSIS. 

KC1O3  +  6HC1  -»  KC1  +  3H20  +  6CI 

If  HN03  or  H2S04  be  substituted  for  HC1,  chlorin  dioxid  is  the  main 
product  and  an  explosion  results,  especially  if  the  acid  be  concentrated 
and  the  mixture  heated.  Hence  nothing  larger  than  a  very  minute  crys- 
tal of  a  chlorate  should  ever  be  heated  with  concentrated  H2S04. 

The  salts  of  chloric  acid  are  the  salts  of  a  very  strong  acid.  They  are 
considerably  more  stable  than  hypochlorites  and  are  not  decomposed  by 
exposure  to  the  atmosphere  or  on  gentle  warming.  On  heating,  however, 
a  general  reaction  of  all  chlorates  takes  place,  resulting  in  the  formation 
of  perchlorates  and  chlorids  and  the  liberation  of  one-third  of  their  oxy- 
gen content. 

2KC103  ->  KC104  +  KC1  +  02 

Since  perchlorates,  on  further  heating,  break  down  into  chlorids  ac- 
cording to  the  reaction  given  on  page  207,  oxygen  in  large  proportion 
is  liberated  in  this  reaction.  The  presence  of  small  amounts  of  Pb02, 
Mn02,  Fe203,  CuO  or  certain  other  finely  divided  substances  causes  the 
gas  to  be  liberated  more  rapidly  and  at  a  lower  temperature. 

All  chlorates  are  good  oxidizing  agents  because  of  this  decomposition 
and  they  are  used  especially  in  high-temperature  reactions  performed  in 
the  ' '  dry  way. ' '  Matches,  ' '  fire- works ' '  and  certain  explosives  consist  of 
mixtures  of  chlorates  with  some  combustible  substance.  In  the  presence 
of  reducing  agents,  the  formation  of  perchlorates  does  not  take  place, 
hence  each  molecule  of  chlorate  liberates  three  atoms  of  oxygen. 

Under  the  influence  of  strong  reducing  agents,  chlorates  are  reduced 
to  chlorids  either  in  acid  or  alkaline  solution,  and  in  this  way  chlorates 
may  be  differentiated  from  the  more  stable  perchlorates  which  are  not 
broken  down  by  the  ordinary  reducing  agents. 

Chloric  acid  acts  as  a  very  strong  monobasic  acid  and,  like  nitric 
acid,  which  it  resembles  in  many  of  its  properties,  it  colors  a  brucin  solu- 
tion red. 

All  chlorates  are  soluble,  the  potassium  salt,  requiring  21  parts  of 
water  for  its  solution,  dissolving  with  the  least  readiness. 

All  chlorates  are  transposed  by  boiling  with  Na2C03. 

The  barium  ion,  Ba,  added  to  a  solution  containing  the  chlorate  ion, 
CK)4,  gives  no  precipitate,  for  BaC103  is  readily  soluble. 

The  silver  ion,  Ag,  causes  no  precipitation,  for  AgC103  is  soluble. 

Perchloric  Acid,  HC1O4. 
Perchloric  acid  is  a  colorless,  fuming,  oily  liquid,  by  far  the  most 


PEKCHLOKATES.  207 

stable  of  any  of  the  oxygen  acids  of  chlorin ;  yet  it  is  little  known  in  the 
pure  state  because  of  its  tendency  to  decompose  explosively.  With 
water,  it  unites  to  form  a  crystalline  hydrate,  HC104.H2O  which  decom- 
poses at  110°.  A  second  hydrate,  HC104.2H20,  distills  undecomposed 
at  203°  and,  since  it  is  perfectly  stable,  this  hydrate  is  the  form  in  which 
perchloric  acid  is  best  known. 

Perchloric  acid  is  one  of  the  very  strongest  acids  and  in  aqueous  solu- 
tion is  almost  completely  dissociated  into  its  ions.  In  solution,  per- 
chlorates  show  but  little  tendency  to  break  down,  even  when  subjected 
to  a  great  variety  of  conditions.  Although  perchlorates  contain  a  greater 
percentage  of  oxygen  than  the  salts  of  any  other  chlorin  acids,  they  are 
not  by  any  means  good  oxidizing  agents ;  for,  because  of  their  stability, 
they  do  not  readily  give  up  any  of  their  oxygen.  Hence,  perchlorates 
possess  but  slight  bleaching  properties  even  in  boiling  acid  solutions  and 
ordinary  reducing  agents,  such  as  zinc  dust,  Devarda's  alloy,  sulfurous 
acid  and  stannous  or  ferrous  chlorids  are  without  effect.  Since  reactions 
in  solution  are  mainly  the  reactions  of  ions,  the  stability  of  perchlorates 
in  aqueous  solution  is  to  be  attributed  to  the  stability  of  the  perchlorate 
ion. 

The  undissociated  acid  and  salts  are  not  by  any  means  as  stable, 
hence  the  danger  of  explosion  in  attempting  to  handle  pure  perchloric 
acid.  This  may  be  explained,  partially,  by  the  fact  that  perchloric  an- 
hydrid,  C1207,  a  colorless  oily  liquid — B.  P.  82° — is  liable  to  spontaneous 
explosion. 

Ignition  of  perchlorates  decomposes  them  into  chlorids  and  liberates 
their  full  oxygen  content. 

0 

// 
K— 0— 01— 0    -»    K— 01  +  40 

\ 
0 

The  resistance  to  reducing  agents  offered  by  perchlorates  in  solution 
and  their  ready  decomposition  when  the  dry  substance  is  heated,  serves 
to  distinguish  perchlorates  from  salts  of  all  other  chlorin  acids.  For, 
on  treating  a  solution  with  sulfurous  acid  or  some  similar  body,  hypo- 
chlorites  and  chlorates  are  reduced  to  chlorids  and  may  be  removed  from 
solution  by  precipitation  with  silver  nitrate.  If  the  filtered  solution  be 
evaporated,  the  residue  ignited  and  the  aqueous  extract  again  treated 
with  AgN03,  any  further  precipitation  will  be  due  to  the  decomposition 
of  perchlorates.  (See  page  306.) 


208  QUALITATIVE    ANALYSIS. 

All  perchlorates  are  soluble,  the  potassium  salt,  requiring  about  100 
parts  of  water  for  its  solution,  being  the  most  insoluble. 

All  perchlorates  are  transposed  by  boiling  with  Na2C03. 

The  barium  ion,  Ba,  added  to  a  solution  containing  the  perchlorate 
ion,  C104  causes  no  precipitation. 

The  silver  ion,  Ag,  gives  no  precipitate. 

BROMIN.    Br  —  80. 

Bromin  is  a  heavy,  dark  brownish-red  liquid,  sharing  with  mercury 
the  distinction  of  being  the  only  elements  existing  normally  in  the  liquid 
state.  It  passes  readily  into  the  gaseous  phase,  however,  when  exposed 
to  the  air  at  ordinary  temperatures,  forming  a  heavy  red  vapor  which 
attacks  the  mucous  membranes  of  the  nose,  eyes  and  throat,  causing 
serious  discomfort.  It  is  soluble  in  about  30  parts  of  water,  forming  a 
solution  known  as  ''bromin  water,"  from  which  all  the  halogen  may  be 
removed  by  boiling.  The  presence  of  a  soluble  bromid  in  solution  renders 
the  bromin  somewhat  more  soluble.  The  increase  in  solubility  is  not 
equal  to  that  caused  by  the  addition  of  iodids  to  an  aqueous  solution 
of  iodin  but  it  is  doubtless  to  be  explained  on  the  same  basis  (see  page 
210).  Bromin  is  so  much  more  soluble  in  carbon  disulfid,  chloroform  and 
ether  than  it  is  in  water  that  it  may  be  removed  from  its  aqueous  solution 
by  shaking  with  any  of  these  liquids.  These  solvents  form  solutions  the 
color  of  which  varies  from  orange-yellow  to  brownish-red  depending  on 
the  concentration  of  the  bromin. 

Bromin  very  closely  resembles  chlorin  in  chemical  properties  but  its 
action  is  less  energetic.  All  of  its  reactions  are  so  closely  analogous  to 
those  of  chlorin  that  a  detailed  description  will  be  unnecessary.  It  unites 
with  most  of  the  elements,  forming  bromids,  of  which  hydrogen  bromid  is 
the  most  important.  It  abstracts  hydrogen  and  liberates  oxygen  only 
very  slowly  from  water,  but  it  attacks  solutions  of  iodids,  setting  free  this 
halogen  and  forming  bromids.  The  reactions  are  explained  by  the  fact 
that  bromin  passes  more  readily  into  the  ionic  condition  than  iodin  or 
oxygen  do.  (See  page  327.) 

Bromin  acts  as  a  very  strong  oxidizing  agent  and  possesses  bleaching 
properties  similar  to  those  of  chlorin,  the  oxygen  equivalent  of  both  halo- 
gens being  the  same.  On  alkalies,  the  action  is  analogous  and  hypobromites 
and  bromates  are  formed  under  identical  conditions.  Inasmuch  as  these 
compounds  are  of  far  less  importance  than  the  corresponding  compounds 
of  chlorin,  since  the  properties  of  the  two  are  so  very  similar,  hypo- 
bromous  and  bromic  acids  will  not  be  considered  in  detail.  No  acids  or 


BKOMIDS.  209 

salts  of  bromin  analogous  to  chlorous  and  perchloric  acids  are  known,  nor 
have  any  oxids  of  bromin  been  discovered. 

Free  bromin  is  usually  detected  by  its  odor,  color,  bleaching  properties 
and  action  on  iodids. 

Hydrobromic  Acid,  HBr. 

Hydrogen  bromid  is  a  colorless  gas  with  sharp  odor  and  acid  taste, 
closely  allied  to  hydrogen  chlorid  in  all  its  properties.  It  fumes  strongly 
in  the  air  and  dissolves  abundantly  in  water,  forming  a  solution  known 
as  hydrobromic  acid.  A  saturated  solution  consists  of  about  80  per  cent. 
HBr  and  20  per  cent.  H20. 

Hydrogen  bromid  is  much  less  stable  than  hydrogen  chlorid,  hence 
it  yields  far  more  readily  to  reagents  generally.  Oxidizing  agents  convert 
it  into  free  bromin  and  water,  especially  in  concentrated  solutions. 
Atmospheric  oxygen  is  sufficient  to  oxidize  it  slowly,  hence  solution  of 
hydrobromic  acid  soon  turn  dark  because  of  the  liberated  bromin.  Be- 
cause of  the  presence  of  the  free  halogen,  the  corrosive  action  of  hydro- 
bromic acid  is  very  great  and  it  will  dissolve  several  metals  not  appre- 
ciably acted  upon  by  hydrochloric  acid. 

If  free  chlorin  be  used  as  an  oxidizing  agent,  as  is  frequently  the 
case  in  testing  for  bromids,  hydrogen  chlorid  is  formed  in  place  of  hy- 
drogen oxid,  i.  e.,  water.  If  a  great  excess  of  chlorin  be  present,  it  may 
unite  with  the  liberated  bromin  to  form  bromin  chlorid,  BrCl,  which, 
although  decomposed  by  water,  may  exist  in  carbon  disulfid  or  chloroform, 
giving  solutions  which  lack  the  red  color  due  to  bromin  and  are  nearly 
colorless. 

Hydrobromic  acid  is  quite  as  completely  dissociated  in  aqueous  solu- 
tion as  is  hydrochloric  acid.  It  acts  as  a  monobasic  acid  and  forms  salts, 
the  solubility  of  which  corresponds  almost  exactly  with  that  of  the 
chlorids. 

Ignition  of  bromids  produces  the  same  effect  as  with  chlorids.  Usu- 
ally the  salt  fuses  unchanged  and  volatilizes  as  the  temperature  is  raised. 

Heated  with  concentrated  H2S04,  all  bromids  are  converted  into  sul- 
fates  and  hydrobromic  acid  is  liberated.  Hot  concentrated  H2S04  pos- 
sesses oxidizing  properties  (see  page  187)  sufficient  to  decompose  some  of 
the  HBr  set  free.  Hence  free  bromin  can  always  be  detected,  together 
with  the  reduction  products  of  H2S04,  generally  S02. 

All  bromids  are  transposed  by  boiling  with  Na2C03  with  the  exception 
of  silver  bromid  which,  like  the  other  halogen  compounds  of  silver,  re- 
quires fusion  with  the  dry  reagent.  ( See  page  203. ) 

15 


210  QUALITATIVE    ANALYSIS. 

•H- 

The  barium  ion,  Ba,  added  to  a  solution  containing  the  bromid  ion, 
Br,  gives  no  precipitate,  for  BaBr2  is  soluble. 

The  silver  ion,  Ag,  precipitates  white  AgBr.     (See  page  87.) 

IODIN.    1  —  127. 

lodin  is  a  dark  gray,  almost  black,  crystalline  solid,  possessing  a 
metallic  lustre.  It  fuses  at  114°  and  boils  at  184°,  forming  a  deep  violet 
vapor.  If  the  temperature  be  maintained  just  below  its  melting-point, 
however,  the  solid  may  be  readily  sublimed.  It  is  quite  insoluble  in 
water,  one  part  of  iodin  requiring  about  5000  parts  of  water  for  solution. 
The  presence  of  a  soluble  iodid  greatly  increases  its  solubility,  this  action 
being  due  to  the  formation  of  a  periodid  which  gives  rise  to  the  ion  I3. 


This  ion  is  stable  only  in  the  presence  of  free  iodin  and  as  fast  as  the 
free  halogen  is  removed  in  any  reaction,  the  periodid  decomposes  into 
the  simple  iodid,  so  that  all  the  free  iodin  originally  put  into  solution  is 
available. 

Iodin  dissolves  very  readily  in  carbon  disulfid,  chloroform,  ether  and 
alcohol  and  may  be  removed  from  aqueous  solutions  by  shaking  with  any 
of  these  solvents,  alcohol  being  excepted  because  it  is  miscible  with  water. 
Carbon  disulfid  and  chloroform  give  rise  to  purple  solutions,  while 
ether  and  alcohol  are  colored  brown. 

Iodin  produces  a  brown  stain  when  brought  into  contact  with  the 
person.  The  application  of  ammonia  to  a  fresh  stain  will  remove  it. 

Iodin  closely  resembles  bromin  and  chlorin  in  chemical  properties,  but 
its  action  is  far  less  energetic  and  temperatures  higher  than  the  ordinary 
are  usually  required  to  make  it  unite  directly  with'  other  elements.  It 
forms  an  unstable  compound  with  hydrogen,  hence  it  does  not  appreciably 
decompose  water  as  the  other  halogens  do. 

Iodin  acts  as  a  very  weak  oxidizing  agent.  Its  oxygen  equivalent 
is  exactly  that  of  chlorin  and  bromin,  but  its  action  is  by  no  means  as 
vigorous.  Many  substances  which  are  oxidized  by  the  latter  elements  do 
not  interact  with  iodin  and  the  oxidization  products  formed  by  iodin, 
when  it  does  react,  are  frequently  not  so  highly  oxidized  as  are  the  reac- 
tion products  when  chlorin  and  bromin  are  used.  (See  page  189.) 

With  starch  in  aqueous  solution,  iodin  produces  a  deep  blue  color 
which  is  due  to  a  body  formed  by  the  addition  of  the  iodin  to  the  starch. 
By  adding  alcohol  or  by  boiling  a  solution  containing  the  starch-iodid 


IOD1N.  211 

compound,  it  is  decomposed  and  the  blue  color  vanishes,  to  reappear  on 
dilution  or  on  cooling,  provided  the  boiling  has  not  been  continued  too- 
long. 

Free  iodin  is  usually  detected  by  the  starch-iodid  reaction,  the  color 
of  its  vapor,  or  of  its  solutions  in  carbon  disulfid  or  chloroform. 

Iodin  acts  on  alkalies  in  a  manner  exactly  analogous  to  chlorin  and 
bromin,  forming  hypoiodites  and  iodates.  Hypoiodites,  however,  are 
more  unstable  than  the  corresponding  salts  of  the  other  halogens  and  they 
have  only  a  temporary  existence  even  in  solution. 

lodic  acid,  HI03,  is  a  white  crystalline  solid,  very  soluble  in  water. 
At  170°  it  loses  water,  forming  the  colorless  crystalline  anhydrid,  I205, 
which  decomposes  into  its  elements  at  300°.  lodic  acid  acts  as  a  very 
strong  monobasic  acid,  forming  salts  which  are  not  as  soluble  as  the 
corresponding  chlorates.  Few  iodates  are  highly  insoluble,  however,  the 
silver  salt,  requiring  about  27,000  parts  of  water  being  the  least  soluble. 

Periodates  cannot  be  formed  by  heating  iodates,  analogous  to  the 
method  for  the  production  of  perchlorates ;  for,  on  ignition,  iodates  break 
down  into  iodids  or  oxids.  Periodic  acid  may  be  formed  by  the  action, 
in  alkaline  solution  or  in  the  ' '  dry  way, ' '  of  very  strong  oxidizing  agents 
on  iodin.  In  its  general  reactions  it  resembles  iodic  acid. 

Hydriodic  Acid,  HI. 

Hydrogen  iodid,  like  the  hydrogen  compounds  of  other  halogens,  is  a 
colorless  gas  with  sharp,  pungent  odor  and  acid  taste.  It  is  abundantly 
soluble  in  water,  the  solution  being  known  as  hydriodic  acid. 

Hydrogen  iodid  is  far  less  stable  than  any  other  halogen  compound 
of  hydrogen.  The  dry  gas  decomposes  very  largely  into  hydrogen  and 
iodin,  and  aqueous  solutions  turn  dark  after  a  very  short  time  because 
of  the  separation  of  iodin.  This  spontaneous  decomposition  into  its 
elements  makes  hydriodic  acid  a  very  strong  reducing  agent,  even  in  di- 
lute solution, — a  property  which  hydrochloric  and  hydrobromic  acids 
show  only  in  concentrated  solutions.  Hence,  in  the  presence  of  sub- 
stances which  can  be  made  to  part  with  a  portion  of  their  oxygen  or 
other  equivalent  element,  hydriodic  acid  reacts  vigorously. 

Ordinary  oxidizing  agents  convert  hydriodic  acid  into  free  iodin  and 
water,  but  vigorous  reagents,  such  as  permanganate  and  chlorin  in  excess, 
oxidize  the  liberated  iodin  to  iodic  acid,  HI03.  In  carbon  disulfid  or  chlo- 
roform colorless  IC1  may  be  formed.  (See  also  page  209.) 

An  idea  of  the  comparative  stability  of  the  hydrogen  compounds  of 
the  halogens  may  be  gained  by  a  consideration  of  the  following  facts. 


212  QUALITATIVE    ANALYSIS. 

Nitrous  acid  liberates  iodin  from  dilute  solutions  of  hydriodic  acid,  but 
does  not  oxidize  hydrochloric  or  hydrobromic  acids  except  in  concentrated 
solutions.  Hydrofluoric  acid  is  not  oxidized  under  any  conditions  by 
nitrous  acid  or  any  other  oxidizing  agent.  Dichromic  acid  in  dilute  solu- 
tions sets  free  iodin;  bromin  is  liberated  if  the  solution  be  somewhat 
more  concentrated,  but  hydrochloric  acid  is  oxidized  only  in  concentrated 
solutions. 

Hydriodic  acid  is  very  completely  dissociated  in  aqueous  solution  and, 
though  very  unstable,  is  quite  as  strong  as  the  other  more  stable  halogen 
acids.  It  is  monobasic  and  forms  salts,  the  solubility  of  which  agrees  very 
closely  with  that  of  the  chlorids,  with  the  exception  that  mercuric  iodid  is 
quite  insoluble. 

Ignition  of  iodids  indicates  the  unstable  character  of  the  acid  from 
which  they  are  derived ;  for,  when  heated  in  the  atmosphere,  they  are  con- 
verted into  the  oxid  of  the  metal  with  the  liberation  of  iodin.  The  salts 
of  the  alkali  metals  are  most  stable  and  may  be  fused  without  appreciable 
decomposition. 

Heated  with  concentrated  H2S04,  all  iodids  are  converted  into  sul- 
f ates,  and  iodin  and  water  are  liberated ;  for  the  hydriodic  acid,  at  first 
set  free,  is  oxidized  by  the  hot  H2S04.  As  reduction  products  of  the 
latter,  sulfur  dioxid,  sulfur  and  hydrogen  sulfid  may  appear. 

All  iodids  are  transposed  by  boiling  with  Na2C03,  except  silver  iodid, 
which  may  be  decomposed  by  fusion  with  the  dry  reagent.  (See  similar 
action  of  chlorids,  page  203.) 

The  barium  ion,  Ba,  added  to  a  solution  containing  the  iodid  ion,  I, 
causes  no  precipitation,  for  BaI2  is  soluble. 

The  siver  ion,  Ag,  precipitates  yellow  Agl.    (See  page  87.) 
The  lead  ion,  Pb,  precipitates  golden-yellow  PbI2.     (See  page  145.) 
The  mercuric  ion,  Hg,  precipitates  vermilion  HgI2.     (See  page  109.) 
The  mercurous  ion,  (Hg)2,  precipitates  olive  (Hgl)2.   (See  page  111.) 

Group  7,  Second  Type. 

Manganese  55. 

Manganese  is  the  only  known  element  of  this  type.  It  stands  in  about 
the  same  relation  to  the  halogen  group  that  chromium  does  to  the  sulfur 
group.  In  the  free  state,  manganese  is  a  typical  metal  and,  while  forming 
compounds  which  show  a  close  similarity  in  chemical  formula,  the  proper- 
ties of  many  of  its  compounds  are  quite  different  from  the  corresponding 
compounds  of  the  halogen  elements.  In  the  divalent  and  trivalent  condi- 


MANGANESE.  213 

tions,  manganese  acts  as  a  weak  base-forming  element;  tetravalent  man- 
ganese is  intermediate  in  character,  like  tetravalent  lead  or  tin,  while 
hexavalent  and  heptavalent  manganese  possess  strong  acid-forming  prop- 
erties. It  is  only  in  this  last  stage  of  oxidation  that  manganese  shows 
similarity  to  the  halogen  group. 

MANGANESE.    Mn  — 55. 

Manganese  is  a  brittle,  grayish-white  metal.  Although  it  is  an  almost 
constant  companion  of  iron  in  natural  compounds,  the  pure  metal  is 
exceedingly  uncommon.  It  is  readily  attacked  by  the  constituents  of  the 
atmosphere,  forming  oxids  or  * '  rusts. ' ' 

Manganese  dissolves  readily,  even  in  acetic  acid,  forming  salts  in 
which  the  metal  is  bivalent  and  acts  as  a  base-forming  element.  During 
the  passage  into  the  ionic  condition,  a  large  amount  of  energy  is  liberated. 
This  causes  the  salts  to  be  exceedingly  difficult  to  reduce  to  the  metal. 
On  the  other  hand,  the  whole  tendency  of  the  element  is  to  pass  into  a 
state  of  higher  oxidation,  becoming  successively  tri-,  tetra-,  hexa-,  and 
finally  hepta-valent.  As  the  valence  increases,  there  is  a  gradual  disap- 
pearance of  the  basic  properties  which  characterize  the  bivalent  condi- 
tion and  a  corresponding  increase  in  acidic  properties  until  the  strongly 
acid  heptavalent  condition  is  reached.  (See  page  45.) 

Five  oxids  of  manganese  are  known,   as  the  following  formulae 

indicate : 

0 

•Kir  rv  Mn=O 

o  Mnr°v  o     i  \ 


/   \  J 

Mn          Mn—  0  Mn     Mn 


I 

I 
° 


\  /  \ 

0 


o 


Mn-°  Mn=0 

'    '        ;  -  > 

Divalent  manganese  oxid,  MnO,  is  the  only  oxid  readily  and  com- 
pletely soluble  in  dilute  acids.  There  is  formed  in  this  way  a  whole  series 
of  salts  in  which  the  metal  is  bivalent  and  acts  as  a  base-forming  element, 
about  as  strong  as  ferrous  iron  or  zinc.  Manganese  salts  are  pale  pink 
when  containing  water  of  crystallization.  In  solution  the  color  is  not  of 
sufficient  intensity  to  be  noticeable.  Anhydrous  salts  are  generally  color- 
less. Hydrolysis  does  not  take  place  in  cold  solutions  to  an  extent  suffi- 


214  QUALITATIVE    ANALYSIS. 

cient  to  cause  precipitation  of  the  hydroxid  when  the  free  acid  in  solution 
is  neutralized  by  the  addition  of  BaC03  or  similar  compounds. 

Trivalent  manganese  is  but  little  known  and  is  unimportant.  Only 
a  few  stable  compounds  are  definitely  known  —  the  sulfate,  Mn2(S04)3, 
and  the  '  '  alum,  '  '  Mn2  (  SOJ  3.K2S04.24H20,  or  better,  KMn  (  S04  )  2.12H,0, 

//° 
besides,  possibly,  the  hydroxid,  Mn  \        .     Other  salts  are  known  only 

OH 

in  concentrated  solutions  of  their  respective  acids  and  are  decomposed 
by  heat  and  dilution.  They  are  of  a  dark  brown  color  (see  page  277).  A 
consideration  of  all  the  known  facts  indicates  that  trivalent  manganese  is 
a  weaker  base-forming  element  than  manganese  in  the  divalent  condition. 
Tetravalent  manganese  is  better  known  and  acts  as  a  very  weak  base- 
forming  as  well  as  weak  acid-forming  element.  This  is  suggested  by  the 
fact  that  the  black  dioxid,  Mn02,  is  not  acted  upon  by  dilute  acids  in  the 
cold.  Fairly  concentrated  acid  is  necessary  to  dissolve  it.  When  so 
treated,  there  is  evidence  of  the  existence  of  salts,  such  as  MnCl4,  in  the 
dark-colored  solution,  but  they  are  less  stable  than  the  trivalent  salts  and, 
on  heating  or  on  dilution,  they  decompose  according  to  the  typical 
reaction  : 


No  salts  are  known  out  of  solution. 

The  partially-dehydrated  hydroxid,  MnO(OH)2,  is  known  as  (meta) 
manganous  acid,  for  it  turns  litmus  red  slowly  and  acts  as  a  weak  acid. 
It  is  but  little  soluble  in  alkalies  because  the  alkali  manganites  are  but 
little  soluble  and,  furthermore,  are  almost  completely  hydrolyzed.  In 
acids  it  dissolves,  forming  the  unstable  salts  already  mentioned,  thus 
showing  that  its  nature  has  been  made  basic  by  the  presence  of  the 
stronger  acids.  If  H2Mn03  be  an  acid  like  H2C03  or  H2Sn03,  then  Mn02 
is  an  acid  anhydrid,  as  the  equation  shows,  and  its  insolubility  in  dilute 
acids  is  thus  explained. 

0—  H  O 


-»    Mn         +  H20 

\  \ 

0—  H  O 

When  Mn304,  like  Pb304  (see  page  143),  is  treated  with  dilute  acids 
two-thirds  of  the  manganese  is  dissolved  while  one-third  is  left  behind  as 
insoluble  MnO2.  It  is  therefore  to  be  regarded  as  the  manganous  salt  of 


MANGANATES.  215 

ortho-manganous  acid   (not  definitely  known),  as  is  indicated  by  the 
formulae  : 

H-0\ 


H—  O\n/r  \0\wr 

H-0/Mn 


H— O/  \0/ 

Ortho-inangauous  acid.  Mangauous  ortho-manganite. 

Hexavalent  manganese  acts  as  an  acid-forming  element  entirely  and 
possesses  no  base-forming  properties.  All  manganese  compounds,  whether 
higher  or  lower  oxidized,  come  into  this  condition  on  heating  with  alkalies, 
forming  intensely  green  salts  of  manganic  acid  which  are  analogous  to 
chromates  and  sulfates  as  the  formulae  indicate: 

K— 0  O  K— O  O  K— O  O 

Mn  Cr  S 

/      \  /     \  /  \ 

K— 0  0  K— 0  O  K— 0  0 

The  properties  of  these  salts  would  indicate  that  H2MnO4  is  a  much 
stronger  acid  than  H2Mn03,  but  not  quite  so  strong  as  chromic  acid. 

The  acid  itself  is  entirely  unknown  for,  as  soon  as  it  is  set  free  from 
one  of  its  salts,  it  decomposes  into  a  higher  and  a  lower  oxidized  product, 
as  the  equation  shows : 

H— 0  0  0      H— 0 

3  Mn        -» 2H—-0— Mn=0  +  Mn=0  +  H20 

/      \  \  / 

H— 0  0  O      H— 0 

The  solution  of  a  manganate  is  green  in  color.  It  contains  the  ion, 
Mn04,  which  is  stable  only  in  the  presence  of  hydroxyl  ions. 

The  manganate  ion,  Mn04,  differs  entirely  in  its  properties  from 
the  permanganate  ion,  Mn04,  just  as  Sn  differs  from  Sn,  and  (Fe  =  C6N6) 
from  (Fe  =  C6N6).  In  the  presence  of  II  ions,  even  the  concentration 
existing  in  pure  water,  Mn04  changes  into  Mn04. 

3Mn04  +  4H  -»  2Mn04  +  H2Mn03  +  H20 

Hence  it  is  that  an  alkaline  solution  of  a  manganate,  by  absorption  of 
C02  from  the  atmosphere,  is  gradually  changed  to  the  pink  color  charac- 
teristic of  permanganates. 


216  QUALITATIVE    ANALYSIS. 

Heptavalent  manganese,  the  highest  state  of  oxidation  that  can  be 
obtained  with  this  element,  is  reached  only  in  acid  solution  by  the  action 
of  strong  oxidizing  agents  on  lower  oxidized  compounds.  In  this  condi- 
tion manganese  is  entirely  acidic  in  nature.  The  oxid,  Mn207,  is  a 
heavy  dark-colored  oil  which  parts  with  its  oxygen  readily  to  form  com- 
pounds of  lower  oxidation.  It  is  the  anhydrid  of  permanganic  acid, 
HMn04,  a  strong  acid,  all  the  compounds  of  which  are  characterized  by 
a  rich  purple  color,  becoming  pink  in  dilute  solution. 

All  compounds  of  manganese,  heated  in  the  oxidizing  flame  before 
the  blowpipe  in  a  bead  of  borax,  or  sodium  metaphosphate,  produce  an 
amethyst  color  due  to  the  formation  of  manganic  oxids  (Mn203  or 
Mn304).  If  manganese  is  present  in  relatively  large  amounts,  the  bead 
is  brown  and  resembles  the  color  given  by  iron  or  nickel.  (See  pages  221 
and  232.) 

Heated  with  oxidizing  agents,  such  as  KC103  or  with  Na2C03  to 
which  a  little  KN03  has  been  added,  all  manganese  compounds  give  a 
green  color  due  to  the  formation  of  manganates  of  the  alkali  metals. 

Boiled  with  HN03  and  Pb02  (or  Pb304),  all  manganese  compounds, 
soluble  under  these  conditions,  give  a  pink  color  to  the  solution,  due  to 
the  formation  of  permanganic  acid.  The  color  is  best  observed  after 
allowing  the  suspended  matter  to  settle. 

The  common  manganese  ions  are  Mn  and  Mn04.  As  the  reactions 
of  these  ions  are  of  importance,  they  will  be  considered  more  in  detail. 

The  Manganous  Ion,  Mn. 

For  the  Mn  ion,  a  solution  of  any  soluble  salt  will  serve.  The  nitrate, 
Mn(NO3)2.6H2O,  is  a  deliquescent,  very  soluble  salt,  melting  in  its  own 
water  of  crystallization  at  temperatures  but  little  above  the  ordinary. 

The  hydroxyl  ion,  OH,  precipitates  white  Mn(OH)2,  insoluble  in 
excess  of  the  precipitant,  thus  indicating  that  bivalent  manganese  pos- 
sesses no  acid-forming  properties.  Unlike  the  hydroxids  of  cobalt,  nickel 
and  zinc,  it  does  not  form  manganese-ammonia  ions  and,  therefore,  does 
not  dissolve  in  NH4OH.  With  ammonium  salts,  however,  it  forms  com- 
plex salts  of  the  type,  (NH4)2(MnCl4),  which  is  characteristic  of  ferrous 
iron,  cobalt,  nickel,  zinc  and  magnesium  as  well.  For  this  reason,  when 
the  ionization  of  NH4OH  is  repressed  by  the  presence  of  NH4C1  in  con- 
siderable quantities,  NH4OH  produces  no  precipitate  of  Mn(OH)2. 

Manganous  hydroxid  slowly  absorbs  oxygen,  forming  brown  man- 
ganous  acid  according  to  the  equation : 


MANGANOUS    SALTS.  217 

H— O 

2Mn(OH)2  +  02    -*    2  Mn=0 

H— 0 

This  change  takes  place  quite  rapidly  in  hot  solutions  containing  the 
OH  ion  and  is  immediate  in  the  presence  of  oxidizing  agents.  Manganous 
acid  reacts  with  manganous  hydroxid  to  form  brown  manganese  manga- 
nite,  which  is  the  oxid  Mn203,  according  to  the  equation : 

0_H      H— 0  0 

/  \  /  \ 

Mn  +  Mn= 0-^Mn          Mn  =  0  +  2H20 

\  /  \  / 

0— H      H— 0  O 

Neither  the  acid  nor  the  manganite  is  soluble  in  NH4C1,  hence,  on 
standing,  a  brown  precipitate  separates  from  hot  solutions  of  manganous 
salts  made  alkaline  with  NH4OH,  even  in  the  presence  of  much  NH4CL 

The  carbonate  ion,  C03,  precipitates  white  MnC03  even  in  the  pres- 
ence of  NH4C1.  In  this  point  manganese  differs  from  cobalt,  nickel,  zinc, 
and  magnesium.  The  carbonate  is  slowly  oxidized  and  hydrolyzed,  thus 
producing  the  brown  precipitate  to  which  reference  has  already  been 
made. 

The  sulfid  ion,  S,  precipitates  from  neutral  and  alkaline  solutions  a 
hydrated  MnS  which  is  flesh-colored.  It  is  exceedingly  soluble,  even  in 
acetic  acid,  and,  for  this  reason,  the  precipitation  by  H2S  from  neutral 
solutions  is  not  complete  because  equilibrium  ensues. 

MnCl2  +  H2S  <-*  MnS  +  2HC1 

For  this  reason,  the  precipitation  is  best  made  in  alkaline  solution  by 
the  use  of  ammonium  sulfid. 

On  boiling  with  a  large  excess  of  (NH4)2S,  the  hydrated  flesh-colored 
MnS  loses  water,  turning  green  as  the  dehydration  proceeds. 

The  sulfate  ion,  S04,  causes  no  precipitate,  as  the  sulfate,  MnS04. 
4H20,  is  soluble  in  about  3  parts  of  H2O.  A  sulfate  containing  7H20 
and  possessing  the  characteristic  form  of  the  " vitriols"  can  be  obtained 
at  low  temperatures  but  it  is  not  the  ordinary  crystallized  salt. 

The  chlorid  ion,  01,  produces  no  precipitate,  for  MnCl2.4H20  is  solu- 
ble in  less  than  its  own  weight  of  water.  The  salt  is  deliquescent  and 
therefore  does  not  keep  well. 


218  QUALITATIVE    ANALYSIS. 

The  cyanid  ion,  CN,  precipitates  white  Mn(CN)2,  soluble  in  dilute 
acids.  It  dissolves  readily  in  an  excess  of  alkaline  cyanid  forming  a  com- 
plex cyanid  of  the  type,  K4(Mn  =  C6N6).  These  manganocyanids  corre- 
spond to  ferrocyanids  and  cobaltocyanids  (see  pages  223  and  230),  and, 
unlike  the  former  but  like  latter,  the  ion  is  quite  unstable  except  in  the 
presence  of  a  considerable  excess  of  alkaline  cyanid.  If  this  excess  be  not 
present,  potassium  manganocyanid  reacts  with  the  manganous  cyanid 
formed  by  decomposition  to  give  a  green  precipitate  of  potassium  man- 
ganese manganocyanid  according  to  the  equation  : 

K4(Mn  =  C6N6)  +Mn(CN)2->K2Mn(Mn  =  C6N6)  +  2KCN. 

The  precipitate  dissolves  on  the  addition  of  more  KCN  reversing  the 
reaction. 

Manganocyanids,  on  exposure  to  the  air,  slowly  oxidize  to  mangani- 
cyanids  according  to  the  equation  : 

2K4(Mn  =  C6N6)  +  0  +  H20  ->2K3(Mn^C6N6)  +  2KOH 
or,  writing  an  ionic  equation  : 

2(Mn  =  C6N6)  +  0  +  H20  -»  2(Mn  =  CCN6)  +  2  OH 

These  compounds  are  no  more  stable  than  the  manganocyanids  and, 
on  acidifying,  are  entirely  decomposed,  the  manganicyanic  acid  first 
formed  breaking  down  as  shown  by  the  equation  : 


2H3(Mn  =  C6N6)  -*  2Mn(CN)2  +  6#C.V  +  (CN)2 

The  ferrocyanid  ion,  (Fe  =  C0N6),  precipitates  white  Mn2(Fe  = 
C6N6),  soluble  in  acids  of  fair  concentration. 

The  ferricyanid  ion,  (Fe  =  C6N6),  precipitates  brown  Mn3(Fe  == 
C6N6)9,  soluble  in  acids  of  fair  concentration. 

The  phosphate  ion,  P04,  precipitates  white  Mn3(P04)2,  soluble  even 
in  acetic  acid. 

The  Permanganate  Ion,  Mn04. 

By  the  action  of  strong  oxidizing  agents  in  the  presence  of  acids  (as 
already  stated,  page  216)  solutions  are  obtained  which  contain  the  per- 
manganate ion,  Mn04,  as  well  as  H  ions.  Permanganic  acid,  HMn04,  is, 
however,  entirely  unknown  out  of  solution,  for  every  attempt  to  isolate  it 
results  only  in  the  anhydrid,  Mn2O7,  or  more  often  its  decomposition 
products. 


PERMANGANATES.  219 

Many  salts  of  permanganic  acid  are  well  known,  however.  They  are 
all  of  a  dark  red  color  with  a  greenish  iridescence.  All  are  readily  solu- 
ble, giving  purple  solutions  which  change  to  pink  as  dilution  increases. 
Inasmuch  as  the  permanganates  of  the  alkali  metals  give  neutral  aqueous 
solutions,  it  is  evident  that  no  hydrolysis  takes  place  and  that  HMn04 
must  be  a  strong  acid. 

The  alkaline  permanganates  are  by  far  the  most  important  salts. 
Potassium  and  ammonium  permanganates  are  soluble  each  in  about  15 
parts  of  H20  and  separate  from  solution  without  water  of  crystallization. 
Sodium  permanganate,  NaMn04.3H20,  is  much  more  soluble  but  cannot 
be  handled  as  well  as  salts  of  other  alkalies. 

Permanganates,  as  a  class,  are  such  highly-oxidized  compounds  that 
they  are  not  very  stable,  but  tend  to  liberate  a  portion  of  their  oxygen 
and  break  down  into  compounds  of  a  lower  stage  of  oxidation.  Heated 
to  240°,  potassium  permanganate,  which  is  the  most  stable  salt,  decom- 
poses into  the  manganate. 

0          K— 0  0  0 

/  \      //  // 

2K— 0— Mn=0    -»  Mn         +     Mn         +  02 

\  /      \  \ 

0          K— 0  0  0 

In  solution,  the  same  change  takes  place  on  long  standing  and  a 
brown  precipitate  of  Mn02,  more  or  less  hydrated,  forms.  By  boiling 
concentrated  solutions  of  permanganates  in  the  presence  of  strong  alka- 
lies, the  reaction  is  hastened  and  the  presence  of  the  manganate  is  indi- 
cated by  the  color  of  the  solution.  This  is  evide^  "y  the  reverse  of  the 
reaction  by  which  permanganates  are  formed.  (See  page  215.) 

The  extreme  ease  with  which  permanganates  liberate  oxygen  make 
them  the  best  of  oxidizing  agents.  Anything  which  can  be  oxidized  at 
all  in  solution  will  be  oxidized  by  permanganates.  In  acid  solutions,  per- 
mangates  give  up  so  much  oxygen  that  they  are  reduced  from  the  hepta- 
valent  to  the  divalent  condition  and  form  manganous  salts  with  the  acid 
present.  From  two  molecules  of  permanganate,  five  atoms  of  nascent 
oxygen  are  liberated  as  the  equation  shows : 

2HMn04  -»  H20  +  2MnO  +  50  = 

In  the  presence  of  H2SO4,  MnO  dissolves  to  form  manganous  sulfate, 
in  HC1,  MnCl2,  etc.  (See  also  page  49.) 

In  alkaline  solution,  only  three  atoms  of  nascent  oxygen  are  obtained 


220  QUALITATIVE    ANALYSIS. 

from  two  molecules  of  permanganate,  the  manganese  changing  from  the 
heptavalent  to  the  tetravalent  condition  only,  and  a  brown  precipitate 
of  Mn02,  more  or  less  hydrated,  is  formed. 

2KMn04  +  H20  -»  2Mn02  +  2KOH  +  30  = 

The  metathetical  reactions  of  permanganates  are  of  little  importance 
since  all  the  salts  are  soluble.  For  this  reason : 

The  barium  ion,  Ba,  added  to  solutions  containing  the  permanganate 
ion,  Mn04,  gives  no  precipitate. 

The  silver  ion,  Ag,  gives  no  precipitate,  except  in  concentrated  solu- 
tions, for  AgMn04  is  soluble  in  about  100  parts  of  H20  at  ordinary 
temperatures. 

Group  8. 

Iron  55.9;  Cobalt  59;  Nickel  58.7; 

Ruthenium  101.7 ;  Khodium  103 ;  Palladium  106.5 ; 

Osmium  191;  Iridium  193;  Platinum  194.8. 

This  group  is  constituted  somewhat  differently  from  the  other  groups 
in  that  its  members  occur  in  triplets  instead  of  singly.  The  free  elements 
are  typically  metallic.  In  combination  they  act  as  weak  base-formers, 
although  the  first  member  of  each  triplet  may  act  as  an  acid-forming 
element  when  highly  oxidized.  What  might  be  called  the  characteristic 
valence  of  eight  is  shown  by  ruthenium  and  osmium  only,  but,  even  with 
these  as  with  the  other  members  of  this  group,  lower  valences  are  more 
often  exhibited.  The  first  triplet,  iron,  cobalt  and  nickel,  are  the  only 
ones  occurring  in  any  quantity.  Platinum  will  be  considered  in  detail, 
also,  because  of  its  importance  in  the  laboratory. 

IRON  (Ferrum).    Fe  — 55.9. 

Metallic  iron,  when  pure,  is  almost  silver-white  in  color  and  melts 
above  1800°.  Commercial  iron  always  contains  impurities  which  darken 
its  color  and  lower  its  melting-point  very  considerably.  It  is  the  most 
magnetic  of  metals.  It  unites  readily  with  oxygen  at  high  temperatures, 
forming  the  magnetic  oxid,  Fe304.  At  ordinary  temperatures  it  is  per- 
manent in  dry  air  but,  in  the  presence  of  moisture  and  C02,  it  "rusts" 
(see  later) . 

Iron  is  very  soluble  in  dilute  acids,  even  in  acetic  and  much 
weaker  acids,  forming  salts  in  which  the  metal  is  bivalent,  or  in  the  fer- 
rous condition.  The  passage  into  the  ionic  condition  is  accompanied  by 
the  liberation  of  energy,  hence  iron  salts  are  not  easily  reduced  to  the 
metal.  Only  in  the  ' '  dry  way ' '  can  this  be  accomplished.  On  the  other 


IKON.  221 

hand,  iron  will  precipitate  silver,  mercury,  bismuth,  copper,  gold  and 
platinum  from  solutions  of  their  salts. 

Two  oxids  are  known  from  which,  by  treatment  with  acids,  two  series 
of  salts  may  be  derived.  On  treatment  with  acids  the  magnetic  oxid, 
Fe304,  forms  two-thirds  ferric  and  one-third  ferrous  salt,  hence  its 
formula  is  written  as  shown  below  : 


\ 
Fe=0  0 

\  \ 

Fe  =  0  O  Fe 

/  / 

Fe=0  O 

/ 

Fe=0 

Ferrous  oxid.  Ferric  oxid.  Magnetic  oxid. 

Ferrous  salts  are  usually  light  green  or  white  in  color,  with  a  slight 
tendency  toward  hydrolysis,  for  bivalent  iron  acts  as  a  stronger  base- 
forming  element  than  many  other  heavy  metals.  The  salts  of  the  ordi- 
nary acids  are  often  soluble  but  the  solutions  become  turbid  because  of 
oxygen  in  the  atmosphere  and  a  precipitate  of  ferric  oxysalts  forms. 

Ferrous  salts  pass  into  the  corresponding  trivalent  ferric  compounds 
on  exposure  to  the  atmosphere  and,  more  rapidly,  under  the  influence 
of  oxidizing  agents.  By  the  action  of  the  stronger  reducing  agents,  the 
reverse  process  is  accomplished  and  ferric  salt  becomes  ferrous. 

Ferric  compounds  are  usually  of  a  red  or  deep  yellow  color,  the  salts 
of  the  common  acids  being  soluble,  hydrolyzing  to  a  very  great  degree  and 
tending  to  form  insoluble  basic  or  oxysalts  which  require  the  presence 
of  free  acid  to  keep  them  in  solution.  Ferric  iron  is  much  less  basic  than 
ferrous  iron  yet  shows  no  acid-forming  properties.  When  the  valence 
rises  to  six,  iron  acts  as  an  acid-forming  element.  The  formula  of  potas- 
sium ferrate,  K2Fe04,  shows  its  analogy  to  K2MnO4,  K2Cr04  and  K2S04. 

Heated  before  the  blowpipe  with  Na2C03  on  charcoal,  all  iron  com- 
pounds are  reduced  to  the  metal.  The  temperature  is  usually  not  high 
enough  to  make  the  metal  sufficiently  liquid  to  collect  in  a  globule  but  a 
black  magnetic  crust  or  fine  gray  particles  are  formed  instead. 

Heated  in  a  borax-bead  before  the  blowpipe  in  the  oxidizing  flame, 
iron  salts  give  a  yellow-to-brown  color  when  hot,  becoming  much  lighter 
when  cold.  In  the  reducing  flame,  the  bead  is  green. 

Of  the  iron  ions  the  ferrous  ion,  Fe,  and  the  ferric  ion,  Fe,  are  the 
most  important.  Besides  these  kathions  iron  is  a  component  of  several 


222  QUALITATIVE    ANALYSIS. 

anions.  The  ferrate  ion,  FeO4,  is  of  little  importance.  Many  complex 
anions  are  known  of  which  the  ferrocyanid  ion,  (Fe  =  C6N6),  and  ferri- 
cyanid  ion,  (Fe=C6N6),  are  the  best  known. 

The  Ferrous  Ion,  Fe. 

For  the  reactions  of  the  Fe  ion,  a  solution  of  the  chlorid  or  sulfate 
are  satisfactory. 

The  hydroxyl  ion,  OH,  precipitates  from  solutions  of  pure  ferrous 
salts  protected  from  all  oxidizing  agents,  white  Fe(OH)2,  readily  soluble 
in  NH4C1,  as  are  the  hydroxids  of  cobalt,  nickel,  manganese  and  zinc. 
Unless  special  precautions  are  taken,  however,  the  precipitate  is  dark 
green  in  color,  owing  to  partial  oxidation  which  may  progress  until  red 
ferric  hydroxid  is  finally  produced.  As  the  oxidation  increases,  the  solu- 
bility in  NH4C1  and  the  similarity  to  the  zinc  group  in  general  decreases. 

The  carbonate  ion,  C03,  in  the  absence  of  air,  precipitates  white 
FeC03,  otherwise  a  green  partially-oxidized  product.  As  the  oxidation 
increases,  hydrolysis  increases  until  finally  ferric  hydroxid  is  precipitated 
and  carbonic  acid  formed  which,  for  the  most  part,  separates  into  CO2 
and  H20. 

Ferrous  carbonate  dissolves  in  water  containing  C02,  forming  the 
acid  carbonate, 

H— 0— C  =  0 

/ 

0 

/ 
Fe 

\ 

O 

\ 
H— 0— C  =  0 

In  this  condition  iron  exists  in  chalybeate  springs  and,  on  exposure  to 
the  air,  is  precipitated  as  ferric  hydroxid. 

The  ' '  rusting  ' '  of  iron  may  be  explained  by  these  same  reactions. 
Carbonic  acid  attacks  the  metal,  liberating  hydrogen  and  forming  ferrous 
acid-carbonate  which  oxidizes  to  ferric  acid-carbonate.  The  latter  com- 
pound hydrolyzes  completely  into  ferric  hydroxid  and  the  original 
amount  of  carbonic  acid.  Ferric  hydroxid,  more  or  less  dehydrated,  is 
"  iron  rust." 

The  sulfid  ion,  S,  precipitates  black  FeS,  readily  soluble  in  acids,  even 
in  acetic  acid,  but  insoluble  in  alkalies  or  their  sulfids.  Because  of  these 


FERROUS    SALTS.  223 

solubility  relations,  H2S  passed  into  an  acid  solution  causes  no  precipi- 
tate. From  a  neutral  solution,  FeS  is  only  partially  precipitated.  An 
equilibrium  ensues  which  may  be  expressed  by  the  equation : 

FeCl2  +  H2S  «-»  FeS  +  2HC1 

The  sulfate  ion,  S04,  produces  no  precipitate,  as  ferrous  sulfate  is 
soluble  in  about  its  own  weight  of  water  at  ordinary  temperatures.  From 
solutions  it  separates  with  7  molecules  of  water  of  crystallization  in  the 
form  which  is  characteristic  of  the  "  vitriols."  As  "  green  vitriol  "  or 
11  copperas,"  it  finds  many  industrial  uses. 

The  chlorid  ion,  Cl,  produces  no  precipitate  as  the  ferrous  halids  are 
all  soluble.  From  aqueous  solution  FeCl2.4H20  separates  and  will  dis- 
solve again  in  less  than  its  own  weight  of  water  at  ordinary  temperatures. 

The  cyanid  ion,  ON,  precipitates  brown  Fe(CN)2,  appreciably  soluble 
only  in  concentrated  acids.  It  dissolves  in  an  excess  of  alkaline  cyanid, 
however,  forming  a  complex  cyanid  of  the  type,  K4(Fe  =  C6N6)  which 
separates  into  K  and  (Fe  =  C6N6)  ions.  The  ferrocyanid  ion  is  one  of 
the  most  stable  of  all  complex  ions,  even  in  very  dilute  solutions  it  decom- 
poses into  the  simple  ions  to  such  a  slight  extent  that  the  concentration 
of  Fe  ions  is  inappreciable  and  no  reactions  of  ferrous  salts  can  be 
obtained.  The  ion  is  stable  under  very  varying  conditions.  On  acidify- 
ing a  solution  of  a  ferrocyanid,  it  is  not  decomposed  but  a  compound  of 

the  composition  H4(Fe  =  C6N6)  may  be  isolated.    In  solution,  this  body 

+ 
ionizes  to   about  the  same  degree   as  sulfuric   acid  into   H  ions   and 

(Fe  =  C6N6)  ions.  Therefore,  it  is  known  as  ferrocyanic  acid,  detailed 
consideration  of  which  is  given  under  a  separate  heading.  ( See  page  130, 
also  127.) 

The  phosphate  ion,  P04,  precipitates  light-green  Fe3(P04)2,  readily 
soluble  in  the  stronger  acids. 

The  ferrocyanid  ion,  (Fe  =  C6N6),  causes  a  white  precipitate  if  all 
oxidizing  agents  be  excluded.  Ordinarily,  the  precipitate  is  light  blue, 
due  to  partial  oxidation.  The  composition  of  the  substance  under  either 
condition  will  vary  according  to  the  relative  amounts  of  the  reagents 
present,  being  Fe2(Fe  =  C6N6)  if  an  excess  of  Fe  ions  be  present,  and 
FeK2(Fe  =  C6N6)  if  an  excess  of  (Fe  =  C6N6)  ions  be  present.  Either 
body  is  insoluble  in  dilute  acids  but  is  decomposed  by  alkalies  with  the 
separation  of  ferrous  hydroxid  and  an  alkaline  ferrocyanid. 

The  ferricyanid  ion,  (Fe==C6N6),  causes  a  dark  blue  precipitate  of 
ferrous  ferricyanid,  Fe3(Fe  =  C6Nc)2,  or  "TurnbuH's  blue."  It  is  in- 


224  QUALITATIVE    ANALYSIS. 

soluble  in  acids  but  is  decomposed  by  boiling  alkalies  into  ferrous  hy- 
droxid  and  alkaline  ferricyanid.  Since  ferricyanids  in  alkaline  solution 
are  energetic  oxidizing  agents,  ferric  hydroxid  is  precipitated  and  alka- 
line ferrocyanid  formed,  in  part,  at  least.  (See  pages  131  and  132.) 

Because  of  the  great  intensity  of  color  produced,  alkaline  ferricyanids 
offer  a  very  delicate  test  for  the  ferrous  iron,  a  test  not  ordinarily  con- 
fused by  the  presence  of  any  other  substance.  Since  ferrous  salts  oxidize 
so  readily,  the  test  should  always  be  made  on  a  fresh  solution  of  the 
original  substance. 

The  thiocyanate  ion,  SON,  is  without  effect  on  a  solution  containing 
only  Fe  ions. 

The  Ferric  Ion,  Fe. 

For  the  reactions  of  the  Fe  ion,  a  solution  of  the  chlorid  serves  best. 

The  hydroxyl  ion,  OH,  precipitates  red  Fe(OH)3  which  loses  water 
in  various  proportions  and  finally  becomes  red  ferric  oxid,  Fe203.  The 
hydroxid,  and  oxid  also  if  formed  by  gentle  heating,  are  readily  soluble 
in  acids,  like  hydroxids  and  oxids  generally,  but,  after  strong  ignition, 
Fe203  is  very  resistant  to  all  ordinary  agents.  By  fusing  with  acid  sul- 
fates,  however,  it  is  changed  to  ferric  sulfate  by  the  nascent  S03  (see 
page  188)  and  so  may  be  obtained  in  solution.  The  hydroxid  shows  no 
acid  properties,  hence  it  is  insoluble  in  alkaline  solutions.  Ferric  hy- 
droxid shows  a  strong  tendency  toward  the  formation  of  pseudo-solutions 
which  are  of  a  deep  red  color. 

In  the  presence  of  tartaric  and  citric  acids,  sugar  and  other  organic 
substances  containing  the  hydroxyl  radicle,  ferric  hydroxid  is  not  precipi- 
tated, owing  to  the  formation  of  complex  bodies  which  do  not  form  Fe 
ions.  Hence  it  is  necessary  that  organic  matter  in  general  be  removed 
from  solution  before  the  attempt  is  made  to  precipitate  ferric  iron. 
Ignition  offers  a  satisfactory  method  of  removing  organic  matter. 

The  carbonate  ion,  C03,  forms,  possibly,  ferric  carbonate,  but,  owing 
to  the  weakness  of  ferric  hydroxid  and  carbonic  acid  as  well,  it  is  com- 
pletely hydrolyzed,  C02  is  liberated  and  Fe(OH)3  is  precipitated. 

The  sulfid  ion,  S,  does  not  precipitate  ferric  sulfid,  for  this  compound, 
formed  only  in  the  ''dry  way,"  is  completely  hydrolyzed  on  treatment 
with  water.  If  H2S  is  passed  into  a  solution  of  FeCl3,  for  instance,  if  fer- 
ric sulfid  formed,  HC1  would  be  formed  as  well.  By  hydrolysis  of  the  sul- 
fid, ferric  hydroxid  would  be  formed  which  would  dissolve  in  the  HC1  pro- 
duced to  give  the  original  salt,  FeCl3.  Hence  no  such  action  is  apparent. 

Hydrogen  sulfid  is  a  reducing  agent,  however,  and,  as  such,  changes 


. 


FEREIC    SALTS.  225 

the  ferric  ions  into  ferrous  ions,  liberating  sulfur  at  the  same  time  as 
the  equations  indicate: 

2FeCl3  +  H2S  ->  2FeCl2  +  2HC1  +  S 
2Fe  +  H2S  ->  2Fe  +  2H  +  S 

In  the  presence  of  acids,  the  color  changes  from  yellow  to  light 
green  and  the  precipitated  sulfur  renders  the  solution  turbid.  Since 
FeS  is  insoluble  in  water  and  alkalies,  if  the  sulfid  ion  is  introduced  in 
excess  into  an  alkaline  solution  of  a  ferric  salt,  the  excess  of  S  ions  will 
react  with  the  Fe  ions,  formed  during  the  reduction,  and  black  FeS  will 
be  precipitated.  -  / 

The  sulfate  ion,  S04,  produces  no  precipitate*;  as  Fe2(S04)3  is  very 
soluble.  It  deliquesces  greatly  under  ordinary  conditions  and  soon  ab- 
sorbs enough  moisture  to  pass  into  solution. 

The  chlorid  ion,  Cl,  causes  no  precipitate,  for  FeCl3  is  very  deliques- 
cent and  very  soluble,  like  other  ferric  halids. 

If  the  normal  salt,  ferric  chlorid  (or  sulfate  or  nitrate)  be  dissolved 
in  water,  the  solution  will  have  an  acid  reaction  owing  to  hydrolysis 
because  of  the  weak  basic  properties  of  trivalent  iron. 

FeCls  +  3H20  <-»Fe(OH)3  +  3HC1 

Because  Fe(OH)3  is  insoluble,  the  reaction  will  continue  in  the  di- 
rection toward  the  right  much  beyond  the  point  at  which  equilibrium 
would  result  if  Fe(OH)3  were  soluble.  Nevertheless,  an  equilibrium  is 
produced,  the  balance  being  brought  about  by  the  opposition  of  the 
two  reactions;  on  the  one  side,  hydrolysis  tending  to  produce  Fe(OH)3, 
on  the  other  side,  the  action  of  HC1  on  Fe(OH)3  to  produce  the  original 
salt.  The  presence  of  free  acid  is  indicated  by  the  reaction  of  the 
solution  to  litmus.  The  presence  of  the  free  base  by  the  turbidity  of  the 
solution.  Since  hydrolysis  increases  with  the  temperature,  the  turbidity 
is  increased  by  boiling. 

If  the  acid  be  removed  in  any  way  as  fast  as  it  is  formed  by 
hydrolysis,  no  equilibrium  will  ensue  and  the  reaction  will  proceed  en- 
tirely in  the  direction  from  left  to  right.  Barium  or  other  insoluble 
carbonate  serves  admirably  to  remove  all  free  acid,  since,  by  the  inter- 
action, carbonic  acid  is  formed  which,  for  the  most  part,  breaks  down 
into  C02  and  H2O.  By  continually  reducing  the  concentration  of  acid, 
BaCO3  causes  complete  hydrolysis  of  the  ferric  salts  of  strong  acids  and 
the  complete  precipitation  of  Fe(OH)3.  Salts  of  aluminum  and  tri- 
valent chromium  act  similarly. 
16 


226  QUALITATIVE    ANALYSIS. 

The  cyanid  ion,  CN,  causes  the  precipitation  of  Fe(OH)3,  ferric 
cyanid  being  unknown.  By  the  union  of  the  hypothetical  Fe(CN)3  with 
3KCN,  a  complex  cyanid  might  theoretically  be  formed,  having  the 
formula,  K8(FeesC6N6).  Such  a  compound,  potassium  ferricyanid,  is 
known  and  may  be  formed  by  the  oxidation  of  potassium  ferrocyanid. 
It  ionizes  into  K  and  (Fe=C6N6)  ions,  the  latter  being  very  stable. 
The  corresponding  hydrogen  compound,  H3(Fe=C6N6)  can  be  isolated. 
It  separates  H  ions  and  is  therefore  known  as  ferricyanic  acid,  under 
which  heading  it  is  treated  more  in  detail.  (See  pages  131,  also  127.) 

The  phosphate  ion,  P04,  precipitates  yellowish-white  FeP04,  which, 
like  phosphates  generally,  is  soluble  in  the  stronger  mineral  acids.  Un- 
like many  phosphates  it  is  insoluble  in  acetic  acid,  and  ferric  compounds 
may  be  separated  from  salts  of  the  metals  included  in  Groups  III  and 
IV  of  the  scheme  of  qualitative  analysis  and  from  magnesium  as 
well,  by  the  introduction  of  the  P04  ion  into  solutions  acidified  with 
acetic  acid. 

The  acetate  ion,  (C2H302),  added  to  a  solution  causes  no  precipitate, 
for  Fe(C2H302)3  is  soluble.  Being  the  salt  of  a  comparatively  weak 
acid  as  well  as  of  a  weak  base,  ferric  acetate  solutions  hydrolyze  to  a 
great  extent,  ferric  hydroxid  showing  especial  tendency  under  these 
conditions  to  form  pseudo-solutions.  This  fact  accounts  for  the  charac- 
teristic deep  red  color  of  ferric  acetate  solutions.  In  a  dilute  boiling 
solution,  an  equilibrium  is  reached  in  which  approximately  two-thirds 
of  all  the  acetic  acid  present  exists  in  the  uncombined  condition.  The 
resulting  precipitates,  having  the  proportions  indicated  by  the  formula, 
Fe(OH)2(C2H302),  is  known  as  basic  ferric  acetate.  Its  appearance  is 
identical  with  ferric  hydroxid,  but  it  is  not  soluble  in  solutions  contain- 
ing acetic  acid  in  concentration  sufficient  to  dissolve  Fe(OH)3. 

The  ferrocyanid  ion,  (Fe  =  C6NJ,  precipitates  dark  blue  ferric  fer- 
rocyanid, Fe4(Fe  =  C6N6)3,  Prussian  blue.  The  compound  is  not  af- 
fected by  concentrated  acids,  but  is  decomposed  by  boiling  alkalies  with 
the  formation  of  ferric  hydroxid  and  alkaline  ferrocyanid. 

Because  of  the  intensity  of  color  produced,  the  ferrocyanid  ion  offers 
a  good  means  of  detecting  the  ferric  ion,  just  as  the  ferricyanid  ion 
serves  to  detect  the  presence  of  the  ferrous  ion.  The  test  should  be  ap- 
plied to  a  fresh  solution  of  the  original  substance,  as,  otherwise,  the 
ferric  ion  may  be  formed  by  the  oxidation  of  ferrous  compounds  by 
the  air. 

The  ferricyanid  ion,  (Fes=C6N6),  produces  no  precipitate,  for 
Fe(Fe=C6N6)  is  soluble.  The  solution  of  ferric  ferricyanid  is  a  deep 
greenish-brown  color,  however. 


COBALT. 


227 


In  testing  for  the  presence  and  condition  of  iron  with  the  complex 
iron  cyanids,  any  color  except  dark  blue  should  be  disregarded.  A  deep 
blue  color  is  produced  whenever  ferrous  ions  are  brought  together  with 
ferri(c)cyanid  ions  or  ferric  ions  with  ferro(us)cyanid  ions,  but  not 
when  ferrous  ions  and  ferro(us)cyanid  ions  come  together  or  ferric  ions 
meet  with  ferri(c)cyanid  ions,  as  the  diagram  indicates. 

Iron  ions.                              Complex  iron-cyanid  ions. 
FerrOus light   blue FerrOcyanid 


Ferric green-brown 


Ferrlcyanid 


The  thiocyanate  ion,  SON,  produces  no  precipitate,  for  Fe(SCN)3 
is  very  soluble.  Undissociated  ferric  thiocyanate  is  blood-red  in  color 
.and  thus  offers  a  very  delicate  test  for  ferric  iron.  The  color  is  intensi- 
fied by  an  excess  of  KSCN,  which  dissociates  to  a  greater  degree  than 
Fe(SCN)3,  and,  therefore,  causes  the  ionization  of  the  latter  compound 
to  be  repressed,  resulting  in  a  larger  quantity  of  the  undissociated  highly- 
colored  salt. 

Ferric  thiocyanate  is  soluble  in  ether  and  in  amyl  alcohol,  neither  of 
which  solvents  cause  dissociation.  By  shaking  an  aqueous  solution  with 
ether  or  amyl  alcohol,  the  Fe(SCN)3  will  be  collected  from  the  water 
and  the  color  intensified.  (See  also  pages  129  and  231.) 

COBALT.     CO  —  59. 

Cobalt  is  a  comparatively  rare,  magnetic,  malleable  metal  of  a  gray 
color  with  a  slight  reddish  tint.  In  many  of  its  properties  and  in  the 
character  of  its  compounds,  it  closely  resembles  iron  and  nickel. 

The  metal  dissolves  slowly  in  HC1  and  H2S04  and  readily  in  HN03, 
forming  compounds  in  which  the  element  is  divalent.  A  small  amount 
of  energy  only  is  liberated  in  passing  into  the  ionic  condition.  For  this 
reason,  the  metal  is  precipitated  from  solutions  of  its  salts  by  zinc,  mag- 


228  QUALITATIVE    ANALYSIS. 

nesium,  and  other  elements  which  liberate  greater  amounts  of  energy  as 
they  become  ions,  i.  e,,  pass  more  readily  into  the  ionic  condition. 

Cobalt  salts  are  of  two  types.  The  cobaltous  series,  in  which  the 
metal  is  divalent,  may  be  derived  from  the  oxid,  CoO.  The  cobaltic  salts, 
in  which  the  element  is  trivalent,  may  be  referred  to  the  oxid,  Co2O3, 
although,  on  treatment  with  acids,  oxygen  (or  its  equivalent)  is  set  free 
and  cobaltous  salts  only  are  formed. 

Co— 0  H— 0  O  00 

\  \  //  /  \  // 

0     +     2  S         -»    2Co  S         +  2H20  +  0 

/  /\  \   /  \ 

Co— 0  H— O          0  00 

A  third  compound  with  oxygen,  Co304,  corresponds  to  Fe304  and  is 
the  most  stable  oxid.  It  is  formed  when  either  of  the  other  oxids  are 
heated.  On  treatment  with  acids,  it  acts  as  does  cobaltic  oxid  and  no 
salts  corresponding  to  it  are  known. 

Cobaltous  salts  are  numerous,  well  known  and  *  generally  stable. 
They  are  of  a  yellow  or  blue  color  when  anhydrous  but  aqueous  solutions 
and  crystals  containing  water  of  crystallization  are  pink.  Most  salts  of 
the  commoner  acids  are  soluble.  They  hydrolyze  but  little  and  resemble 
very  closely  the  corresponding  ferrous  salts.  Unlike  ferric  salts  which 
hydrolyze  greatly,  cobaltous  salts  are  not  precipitated  as  the  hydroxid  by 
treating  their  cold  solutions  with  BaC03. 

Cobaltic  compounds  are  generally  unstable  under  ordinary  condi- 
tions. Aside  from  the  oxid  and  hydroxid,  no  simple  salts  are  definitely 
known.  Possibly  some  of  the  halids  exist  in  solution.  Several  complex 
salts  of  trivalent  cobalt  are  stable  and  well  recognized,  however.  They 
will  be  considered  in  connection  with  the  cobaltous  compounds  from 
which  they  may  be  derived. 

Heated  before  the  blowpipe  with  Na2C03  on  charcoal,  all  cobalt  com- 
pounds are  reduced,  leaving  a  dark  magnetic  crust. 

Heated  in  the  borax-bead,  all  cobalt  compounds  produce  a  deep  blue 
color,  probably  due  to  the  formation  of  cobalt  meta-borate,  Co{B02)2. 

The  common  cobalt  ion  is  Co.  The  trivalent  ion,  Co,  is  not  definitely 
known.  Both  divalent  and  trivalent  cobalt  form  large  numbers  of  com- 
plex salts  of  varying  composition,  most  of  which  give  easily  recognized 
complex  ions.  Complex  cobalt-ammonia  salts  of  two  types  are  known,  one 
of  which  contains  four  molecules  of  NH3  to  one  of  the  metal,  the  other 
6NH3.  The  salts  give  rise  to  the  following  complex  ions : 

Co.(NH3)4  C+o.(NH3)6  Co.(NH3)6 


COBALT    COMPOUNDS.  229 

The  NH3  in  these  compounds  may  be  replaced  more  or  less  completely 
by  H20  and  other  molecules  •  by  CN,  N02,  and  other  radicles ;  and  by  Cl 
and  other  atoms,  giving  rise  to  a  great  variety  of  compounds.  If  the 
substituting  radicle  or  element  is  acidic Jn  nature,  i.  e.,  ordinarily  forms 
ions  carrying  negative  charges,  as  Cl,  CN  and  N02,  the  character  of  the 
complex  ion  changes  gradually,  one  positive  charge  being  removed  or  a 
minus  charge  being  added  as  each  successive  substitution  occurs,  as  the 
following  series  indicates : 

Co".(NH3)6,    Co.(NH3)5.N02,    Co.(NH3)4.2N02,     Co.(NH3)3.3N02, 
Co.  (NH3)  2.4N02,     Co.  (NH.)  .5N02,     Co.6NOa. 

The  last  of  the  above  series — the  cobalti-nitrite  ion — as  well  as  the 
corresponding  cobalticyanid  ion,  Co.GCN,  are  well  known.  By  an  ex- 
tension of  this  principle,  all  the  known  ammonia  compounds  (and  their 

derivatives)  of  any  of  the  metals  may  be  included  in  this  classification. 

++ 
For  the  reactions  of  the  Co  ion  any  soluble  salt  may  be  used.     The 

nitrate,  Co(N03)2.6H20  is  a  red,  deliquescent,  very  soluble  salt,  melting 
in  its  own  water  of  crystallization  at  temperatures  not  much  above  the 
ordinary.  The  acetate,  Co(C2H302)2.4H20  effloresces  and,  therefore, 
is  more  readily  preserved  than  other  cobalt  salts. 

The  hydroxyl  ion,  OH,  added  to  a  solution  containing  the  cobalt  ion, 
Co,  precipitates  a  blue  basic  salt  at  first,  Co(OH)(N03)  for  instance, 
which,  in  the  presence  of  more  OH  ions,  changes  into  the  pink  Co(OH)2. 
A  high  concentration  of  OH  ions  causes  some  cobaltous  hydroxid  to  dis- 
solve, thus  showing  that  it  may  be  forced  to  display  slight  acidic  proper- 
ties. In  NH4OH  and  NH4C1  it  dissolves  readily,  forming  complex 
ammonia  compounds.  In  the  presence  of  sufficient  NH4C1,  therefore, 
NH4OH  produces  no  precipitate  owing  to  the  formation  of  these  com- 
pounds and  to  the  repression  of  its  ionization  as  well. 

On  standing  in  the  air,  more  rapidly  under  the  influence  of  oxidizing 
agents,  especially  of  alkaline  hypochlorites  or  hypobromites  or  of  hydro- 
gen peroxid,  cobaltous  hydroxid  oxidizes  to  brownish-black  cobaltic 
hydroxid,  Co(OH)3. 

The  carbonate  ion,  C03,  causes  a  pink  precipitate  consisting  of  the 
carbonate  and  hydroxid  in  varying  proportions.  This  precipitate  does 
not  fall  in  the  presence  of  ammonium  chlorid  if  (NH4)2C03  is  used  as 
the  precipitating  reagent. 

The  sulfid  ion,  8,  causes  no  precipitate  in  solutions  containing  an 
appreciable  amount  of  acid.  In  neutral,  alkaline  and  in  acetic  acid  solu- 


230  QUALITATIVE    ANALYSIS. 

tions,  provided  they  be  barely  acid  in  reaction,  black  CoS  is  precipitated. 
Although  the  precipitation  is  prevented  by  a  very  small  concentration 

+ 

of  H  ions,  CoS,  once  precipitated,  does  not  dissolve  readily  in  anything 
except  aqua  regia.  This  phenomenon  which  is  exhibited  by  NiS,  also, 
is  probably  due  to  the  existence  of  a  highly  insoluble  modification  into 
which  the  sulfids  first  precipitated  immediately  change.  Although  not 
readily  or  completely  soluble  in  anything  except  aqua  regia,  the  pre- 
cipitated sulfid  does  dissolve  somewhat  in  most  of  the  dilute  acids  even  in 
acetic  acid  and,  when  moist,  is  slowly  oxidized  by  the  air  into  CoS04. 

The  sulfate  ion,  S04,  causes  no  precipitation,  for  CoS04  is  readily 
soluble.  From  solution  it  separates  with  7H2O  in  the  crystal  form  char- 
acteristic of  the  '  '  vitriols.  '  ' 

The  chlorid  ion,  Cl,  causes  no  precipitate  to  fall,  for  CoCl2.6H20  is 
soluble  in  2-3  parts  of  H20. 

The  cyanid  ion,  ON,  precipitates  reddish-brown  Co(CN)2,  insoluble 
in  acids  but  readily  soluble  in  alkaline  cyanids,  forming  a  complex  salt 
of  the  type  K4(Co  =  C6N6),  similar  to  potassium  ferrocyanid.  The 
complex  cobaltocyanid  ion,  (Co  =  C6N6),  is  far  less  stable  than  the  ferro- 
cyanid ion  and,  in  acid  solutions,  is  entirely  broken  up  into  the  simple 
ions.  Consequently  on  acidifying  a  solution  of  a  cobaltocyanid,  Co(CN)2 
is  precipitated  and  HCN  set  free. 

On  treating  with  oxidizing  agents,  cobaltocyanids  are  oxidized  to 
cobalticyanids  according  to  the  reactions  : 


2K4(CO  =  C6N6)  +  0  +  H20-^  2K3(Co  =  C6N6)  +  2KOH 

CN          2K      2  OH 


which  simplifies  to 

C6N6)  +  0  +  H20-»2(Co=C6Ne)+20H 


Cobalticyanids  are  very  similar  to  ferricyanids.  The  ion, 
(Co  =  C6N6),  is  much  more  stable  than  the  ion,  (Co  =  C6N6).  It  is  not 
decomposed  in  acid  solutions  and,  on  acidifying  a  solution  of  a  cobalti- 
cyanid,  neither  HCN  nor  cobaltic  cyanid  are  formed.  From  the  solution 
a  compound  may  be  isolated  having  the  composition  indicated  by  the 
formula,  H3(Co  =  C6N6),  which  separates  H  ions  and,  therefore,  is  to 
be  considered  as  cobalticyanic  acid.  The  alkali  salts  are  soluble  and 
give  with  the  ions  of  the  heavy  metals  characteristic  precipitates.  The 
compounds  of  cobalticyanic  acid  are  not  of  importance,  however,  and 
will  not  be  further  considered. 


COBALT    SALTS.  231 

The  nitrite  ion,  N02,  produces  no  precipitate  in  dilute  solution,  for 
cobaltous  nitrite  is  soluble.  In  concentrated  solutions,  acidified  with 
acetic  acid  (but  not  with  any  mineral  acid),  the  nitrous  acid  formed  oxi- 
dizes Co(N02)2  to  Co(N02)3,  which  will  then  unite  with  KN02  (but  not 
NaN02)  to  form  a  complex  salt,  K3(Co=  (N02)6),  which  is  yellow  and 
fairly  insoluble.  Any  salt  of  nickel  formed  under  the  same  conditions 
is  readily  soluble,  hence  this  reaction  offers  a  means  of  separating  these 
two  metals.  This  separation  requires  several  hours'  standing,  however, 
and  should  be  repeated  once  again  on  the  nitrate  if  a  considerable  quan- 
tity of  cobalt  is  present. 

The  ferrocyanid  ion,  (Fe  =  C6N6),  causes  a  drab  precipitate  of 
very  varying  composition,  approximating  more  or  less  closely  to 
Co2(Fe  =  C6N6).  It  is  insoluble  in  dilute  acids,  but  soluble  in  NH4OH. 

The  ferricyanid  ion,  (Fe  =  C6N6),  causes  a  slimy  reddish  precipitate, 
Co3(Fe  =  C6N6)2,  insoluble  in  dilute  acids  and  in  NH4OH. 

The  thiocyanate  ion,  SON,  causes  no  precipitate,  for  Co(SCN)2  is 
very  soluble.  In  concentrated  solutions  of  a  thiocyanate,  however,  com- 
plex salts  are  formed  which  have  a  beautiful  blue  color.  They  are 
soluble  and  more  stable  in  amyl  alcohol  than  in  water.  In  the  latter 
solvent,  on  dilution,  the  complex  salt  is  decomposed  into  the  simple  salts. 
Inasmuch  as  nickel  does  not  form  a  corresponding  compound,  a  means 
of  detecting  traces  of  cobalt  in  the  presence  of  large  quantities  of  nickel 
is  here  presented.  To  the  solution  to  be  tested  5-10  grams  of  solid 
NH4SCN  is  added  and  the  solution  then  thoroughly  shaken  with  5-10 
ccm.  of  amyl  alcohol.  If  cobalt  is  present,  the  alcohol  layer  will  be 
colored  blue. 

If  ferric  iron  is  present,  Fe(SCN)3  (see  page  227)  will  form,  the 
blood-red  color  of  which  may  mask  the  test  since  ferric  thiocyanate  is 
also  soluble  in  amyl  alcohol.  On  shaking  the  alcohol  layer  with  a  solution 
of  Na2C03,  the  iron  will  be  precipitated  while  the  cobalt  will  not  and  the 
blue  color  will  remain  in  the  alcohol. 

The  phosphate  ion,  P04,  precipitates  pink  Co3(POJ2,  soluble  in 
dilute  acids,  even  in  acetic  acid. 

NICKEL.    Ni  — 58.7. 

Nickel  is  a  malleable,  slightly  magnetic  metal.  It  is  almost  silver- 
white  in  color  and  is  stable  in  the  air.  It  closely  resembles  cobalt  in  the 
character  of  its  compounds. 

The  metal  dissolves  slowly  in  HC1  and  H2S04  and  readily  in  HN03, 
forming  compounds  in  which  the  element  is  divalent.  The  passage  into 


232  QUALITATIVE    ANALYSIS. 

the  ionic  condition  is  accompanied  by  the  liberation  of  a  small  amount 
of  energy  and  the  metal  may  be  precipitated  from  solutions  of  its  salts  by 
metallic  zinc. 

Two  oxids  of  nickel  are  known,  the  green  divalent  nickelous  oxid, 
NiO,  and  the  black  trivalent  nickelic  oxid,  Ni203. 

All  nickel  salts  are  formed  from  the  lower  oxid.  The  series  is  charac- 
terized by  a  green  color  in  solution  and  when  containing  water  of  crystal- 
lization, the  color  changing  to  yellow  as  the  substances  become  anhydrous. 
Many  of  the  salts  of  the  commoner  acids  are  soluble  and  in  solution 
hydrolyze  but  little,  showing  that  nickel,  like  divalent  metals  in  general, 
acts  as  a  stronger  base-forming  element  than  trivalent  metals.  For  this 
reason,  nickel  cannot  be  precipitated  as  the  hydroxid  by  adding  BaC03 
to  cold  solutions  of  its  salts. 

No  trivalent  nickel  compounds  are  known  except  the  hydroxid, 
Ni(OH),. 

Heated  before  the  blowpipe  with  Na2C03  on  charcoal,  nickel  com- 
pounds are  reduced  to  the  metal,  forming  a  dark-gray  crust  which  is 
slightly  magnetic. 

Heated  in  a  borax-bead  in  the  oxidizing  flame,  nickel  produces  a  yel- 
lowish-brown color  very  like  that  given  by  iron.  In  the  reducing  flame, 
the  bead  turns  gray  and  becomes  opaque,  owing  to  the  presence  of  fine 
particles  of  the  metal.  The  borax-bead  can  hardly  be  considered  as  a 
decisive  test  for  nickel,  however. 

The  nickel  ion,  Ni,  is  best  known.  Complex  ammonia  and  cyanogen 
ions  are  recognized. 

For  the  reactions  of  the  Ni  ion,  any  soluble  simple  salt  will  serve. 
The  nitrate,  Ni(NO3)2.6H20,  is  soluble  in  about  2  parts  of  H20. 

The  hydroxyl  ion,  OH,  precipitates  apple-green  Ni(OH)2,  insoluble 
in  an  excess,  thus  indicating  that  nickel  does  not  show  acid  properties. 
In  NH4OH  and  in  NH4C1  it  is  readily  soluble,  forming  a  number  of  com- 
plex ammonia  compounds  similar  to  some  of  the  cobalt  compounds.  For 
this  reason,  NH4OH  (or  (NHJ2C03)  causes  no  precipitate  in  the  pres- 
ence of  NH4C1. 

In  the  presence  of  oxidizing  agents,  especially  alkaline  hypochlorites 
and  hypobromites  (but  not  of  hydrogen  peroxid — distinction  from  cobalt, 
see  page  229),  black  Ni(OH)3  is  formed.  In  alkaline  solutions,  fer- 
ricyanids  act  as  oxidizing  agents  and  produce  the  same  product  on  boil- 
ing. Nickelic  hydroxid  dissolves  readily  in  acids,  forming  divalent 
nickel  salts  and  liberating  oxygen  or  its  equivalent.  (See  page  228.) 


NICKEL    SALTS.  233 

0— H        H— 0  O  0  O 

/  \  //  /   \  // 

2Ni— 0— H  +  2  S        ->2Ni  S        +5IL.O  +  0 

\  /  \  \   /  \ 

0— H        H— 0  0  00 

The  carbonate  ion,  C03,  causes  an  apple-green  precipitate,  varying 
in  composition,  consisting  of  a  mixture  of  Ni(OH)2  and  NiC03. 

The  sulfid  ion,  S,  precipitates  black  NiS  from  solutions  neutral,  alka- 
line or  barely  acid  with  acetic  acid.  In  the  presence  of  a  higher  concen- 

-}- 

tration  of  H  ions  and  of  salts  of  tartaric  and  other  organic  acids,  no 
precipitate  falls,  the  solution  merely  turning  dark.  An  excess  of 
(NH4)2S  exerts  a  similar  action,  precipitating  part  only  of  the  NiS  and 
giving  a  brown  nitrate.  When  once  precipitated,  however,  NiS  is  not 
appreciably  soluble  in  acids  or  in  any  of  these  other  reagents  but  shows 
the  same  phenomena  as  CoS.  (See  page  229.) 

The  sulfate  ion,  S04,  causes  no  precipitate,  as  nickel  sulfate  is  sol- 
uble in  2-3  parts  of  H20.  From  solution  it  crystallizes  with  7H20  in 
the  crystal  form  characteristic  of  the  "vitriols." 

The  chlorid  ion,  Cl,  causes  no  precipitate,  as  NiCl2.6H20  is  soluble 
in  less  than  two  parts j)f  H20. 

The  cyanid  ion,  CN,  precipitates  green  Ni(CN)2,  insoluble  in  dilute 
acids.  In  an  excess  of  an  alkaline  cyanid,  it  is  readily  soluble,  forming 
a  complex  cyanid  of  the  type,  K2Ni(CN)4,  which  forms  the  complex 
ion,  (NiC4N4).  In  the  presence  of  H  ions,  it  is  unstable  and  decomposes 
entirely  into  the  simple  ions,  Ni  and  CN.  By  treating  the  complex 
nickel  cyanid  with  bromin  or  other  oxidizing  agents,  it  is  not  oxidized 
to  a  complex  salt  of  trivalent  nickel,  as  is  the  case  with  the  corresponding 
cobalt  salt,  but  it  is  decomposed  according  to  the  equation : 

K2Ni  ( CN)  4  +  8Br  ^  2KBr  +  NiBr2  +  4CNBr 

In  the  presence  of  NaOH  or  KOH,  Ni(OH)2  is  formed  and  is  oxidized 
by  the  bromin  to  Ni(OH)3,  which  is  precipitated. 

2Ni(OH)2  +  2Br  +  2NaOH  -»  2Ni(OH)3  +  2NaBr 

By  this  means  a  separation  of  cobalt  and  nickel  may  be  effected, 
since,  under  similar  treatment,  a  soluble  cobalticyanid  is  formed.  (See 
page  230.) 

The     ferrocyanid     ion,      (Fe  =  C^N6),     precipitates     light     green 


234  QUALITATIVE    ANALYSIS. 

Ni2(Fe  =  C6N6),  insoluble  in  dilute  acids  but  soluble  in  NH4OH,  forming 
complex  nickel-ammonia  compounds. 

The  ferricyanid  ion,  (Fe  =  C6N6),  precipitates  .  olive-brown 
Ni3(Fe  =  C6N6)2,  insoluble  in  dilute  acid,  but  soluble  in  NH4OH,  form- 
ing complex  nickel -ammonia  compounds.  Since  cobalt  ferricyanid  is 
insoluble  in  NH4OH,  this  offers  a  most  satisfactory  method  of  separating 
nickel  from  cobalt.  (See  page  231,  also  268.) 

The  thiocyanate  ion,  SON,  causes  no  precipitate  and  gives  no  color 
to  the  solution. 

The  phosphate  ion,  P04,  precipitates  apple-green  Ni3(POJ2,  soluble 
in  acids,  even  in  dilute  acetic  acid. 

PLATINUM.    Pt  — 194.8. 

Platinum  is  a  heavy,  grayish- white,  malleable  metal,  melting  at  1770°, 
a  temperature  somewhat  higher  than  the  flame  of  an  ordinary  Bunsen 
burner.  It  resembles  gold  in  many  of  its  properties.  Like  the  latter 
element,  it  doe's  not  tarnish  in  the  atmosphere  at  any  temperature  and 
does  not  dissolve  in  any  acid.  Since  it  is  readily  acted  upon  by  the  halo- 
gens, it  does  dissolve  in  aqua  regia,  forming  hydrochlorplatinic  acid, 
H2PtCle.  Platinic  chlorid,  PtCl4,  is  undoubtedly  a  transition  product 
(see  later). 

The  finely  divided  metal  obtained  by  the  reduction  of  its  salts — 
"platinum  black" — or  by  their  decomposition  by  heat — "platinum 
sponge" — has  the  power  of  absorbing  large  quantities  of  hydrogen,  oxy- 
gen and  other  gases.  Since  these  substances  are  liberated  in  a  very  much 
more  reactive  (nascent)  condition  (see  page  40),  the  presence  of  finely 
divided  platinum  is  a  very  great  aid  in  certain  chemical  reactions. 

Platinum  does  not  pass  voluntarily  into  the  ionic  condition  but  re- 
quires that  energy  be  expended  to  make  it  assume  the  electric  charge. 
Like  gold  it  stands  very  low  in  the  Potential  Series,  hence  from  solutions 
of  platinum  salts,  the  element  is  precipitated  by  the  introduction  of  any 
of  the  other  common  metals. 

With  oxygen,  platinum  does  not  unite  directly  but,  by  indirect  meth- 
ods, two  compounds  with  this  element  may  be  obtained. 

Platinum  monoxid,  PtO,  together  with  the  divalent  platinous  salts 
derived  from  it,  are  of  little  importance  and,  in  general,  resemble  the 
corresponding  higher  oxidized  compounds  which  are  generally  of  a  much 
lighter  color. 

The  tetravalent  platinic  salts,  derived  from  the  dioxid,  Pt02,  are  gen- 
erally soluble  and  of  a  yellowish-brown  color.  They  are  more  stable  than 


PLATINUM.  235 

the  corresponding  auric  salts,  as  is  indicated  by  the  facts  that  oxalic  acid 
and  ferrous  salts  cause  no  reduction  at  ordinary  temperatures  in  acid 
solutions,  although  long  boiling  with  ferrous  salts  in  neutral  solutions 
causes  a  precipitation  of  the  metal.  Stannous  chlorid  does  not  precipi- 
tate metallic  platinum  but  causes  the  formation  of  soluble  platinous  salts. 
Hydriodic  acid  accomplishes  a  similar  reduction  to  the  platinous  condi- 
tion but  hydrogen  sulfid  is  without  effect.  (See  page  89.) 

On  ignition,  all  platinum  salts  are  decomposed  into  the  metal  and, 
when  heated  before  the  blowpipe  on  charcoal  with  Na2C03,  they  leave 
behind  a  dry  infusible  gray  powder,  which  shows  metallic  lustre  when 
rubbed  in  a  mortar. 

Simple  platinum  salts  are  only  infrequently  met  with,  owing  to  a 
very  strong  tendency  toward  the  formation  of  complex  bodies  by  the  addi- 
tion of  many  other  substances.  To  illustrate:  Platinic  chlorid,  PtCl4, 

-^[.  [.  

which  separates  into  the  simple  Pt  and  Cl  ions,  is  but  little  known  be- 
cause it  unites  so  readily  with  hydrogen  chlorid  (or,  analogously,  with 
other  chlorids)  to  form  the  compound  hydrochlorplatinic  acid,  H2PtCl6, 
which  separates,  along  with  hydrogen  ions,  the  chlorplatinate  ion,  PtCl6. 

Salts  of  hydrochlorplatinic  acid  are  of  considerable  importance. 
The  potassium  salt,  K2PtCl6,  and  the  ammonium  salt,  (NH4)2PtCl6,  dis- 
solve to  some  extent  in  water  (see  pages  74  and  78)  but  they  are 
highly  insoluble  in  dilute  alcohol.  They  are  frequently  made  use  of  for 
the  separation  of  ammonium  and  potassium  radicles  from  solution. 

The  hydroxyl  ion,  OH,  added  to  an  acid  solution  containing  tetra- 
valent  platinum,  does  not  precipitate  platinic  hydroxid,  but  forms  in- 
stead the  alkali  chlorplatinate.  Platinic  hydroxid,  when  obtained  under 
carefully  prescribed  conditions,  is  found  to  act  either  as  a  base  or  an 
acid,  dissolving  either  in  acids  or  in  bases,  like  gold  hydroxid. 

The  sulfid  ion,  S,  added  to  a  solution  containing  tetravalent  platinum, 
decomposes  the  chlorplatinate  ion,  PtCl6  and  precipitates  black  PtS2, 
which  is  insoluble  in  all  acids.  It  dissolves  in  aqua  regia,  however. 
Platinic  sulfid  unites  with  alkaline  sulfids  to  form  soluble  thioplatinates 
similar  to  thiostannates  (see  page  141).  In  the  scheme  of  qualitative 
analysis,  therefore,  platinum  appears  in  the  tin  group. 

The  Care  of  Platinum  Ware. 

Although  platinum  is  one  of  the  most  resistant  of  metals,  it  is  in- 
correct to  assume  that  it  may  be  subjected  to  the  action  of  all  chemical 
reagents  under  all  conditions  without  injury.  At  ordinary  temperatures 
only  aqua  regia  and  the  halogens  act  on  it,  but,  at  higher  temperatures, 


236  QUALITATIVE   ANALYSIS. 

many  other  substances  have  a  very  appreciable  action.  Alkaline  hy- 
droxids,  sulfids  and  cyanids  in  the  state  of  fusion  all  attack  it,  although 
alkaline  carbonates  and  borates  are  without  effect.  At  high  temperatures, 
iodin,  sulfur,  phosphorus  and  arsenic  react  vigorously,  and  most  metals, 
except  iron,  form  readily  fusible  alloys.  Carbon  is  not  without  action  on 
highly  heated  platinum  and  vessels  of  this  metal  should  never  be  ignited 
over  a  luminous  flame  which  deposits  soot,  i.  e.,  carbon,  on  a  cold  object 
introduced  into  it. 

Not  alone  must  the  heating  of  any  of  these  substances  in  a  platinum 
vessel  be  avoided,  but  care  should  be  exercised  that  none  of  these  sub- 
stances shall  be  formed  by  the  interaction  at  high  temperatures  of  sub- 
stances which  are  themselves  harmless.  Salts  of  the  metals  (especially 
of  silver,  mercury,  lead,  bismuth,  tin,  zinc,  arsenic  and  antimony)  or  sul- 
fates  or  phosphates  should  never  be  heated  with  reducing  matter,  for  the 
reduction  products  may  completely  ruin  any  platinum  vessel  in  which 
the  ignition  takes  place. 

The  physical  condition  of  platinum  is  a  prominent  factor  in  in- 
fluencing the  action  of  reagents  upon  it.  Platinum  ware  should  always 
be  kept  smooth,  clean  and  polished.  Should  it  become  stained  or  tar- 
nished, it  should  be  boiled  with  dilute  nitric  acid,  or  it  may  be  fused  with 
borax  and  afterwards  digested  with  hot  dilute  nitric  acid.  Platinum 
may  be  polished  conveniently  by  rubbing  between  the  fingers  with  fine 
sea-sand. 


PART  III 


ANALYTICAL 


SYSTEMATIC  ANALYSIS. 

The  systematic  analysis  of  any  solid  substance  consists  of  three  parts:  the 
Preliminary  Examination,  as  given  pages  243-247,  the  Basic  Analysis — testing 
for  kathions  or  metallic  elements — as  given  pages  248-273,  the  Acid  Analysis — • 
testing  for  anions  or  acid  radicles — as  given  pages  275-309. 

The  Preliminary  Examination  is  made  on  small  portions  of  the  original 
solid  substance.  It  should  rarely  be  omitted  since  it  takes  but  a  few  minutes 
and,  if  the  substance  is  homogeneous  (i.  e.,  is  not  a  mixture)  and  is  fairly 
simple,  the  results  of  a  preliminary  examination  may  render  testing  unneces- 
sary. 

The  Basic  Analysis  is  the  systematic  examination  to  determine  the  presence 
or  absence  of  ions  of  the  more  commonly  occurring  metals.  The  tests  consist 
entirely  of  reactions  performed  in  the  "  wet  way,"  thus  requiring  that  the 
original  substance  be  obtained  in  solution.  It  is  pulverized,  therefore,  and 
the  action  of  the  following  solvents  is  tried  in  the  order  given. 

Water.  Boil  about  1  gram  of  the  finely  powdered  substance  with  10-20  ccm. 
of  distilled  water.  If  the  substance  is  readily  and  completely  soluble,  proceed 
at  once  to  the  Basic  Analysis.  If  it  does  not  dissolve  apparently,  evaporate 
1-2  ccm.  of  the  water  extract  in  platinum  to  see  whether  anything  has  gone 
into  solution.  If  an  appreciable  residue  is  left  on  the  platinum,  continue  the 
extraction  of  the  original  substance  with  fresh  portions  of  boiling  water  as 
long  as  anything  continues  to  dissolve.  If  the  substance  is  insoluble  in  water, 
try  the  action  of— 

Hydrochloric  Acid.  Warm  the  residue  left  after  extraction  with  water  with 
dilute  HC1.  If  the  substance  dissolves  readily,  continue  the  addition  of  HC1, 
if  necessary,  until  no  residue  is  left  and  treat  the  solution  at  once  as  directed 
under  Filt.  1  of  the  Basic  Analysis;  for  Group  I,  Div.  A  cannot  be  present  if 
the  substance  is  soluble  without  residue  in  HC1.  If  the  substance  is  not  readily 
soluble  in  the  dilute  acid,  try  the  action  of  concentrated  HC1  and  evaporate  a 
portion  of  the  liquid  in  platinum  to  see  if  anything  dissolves.  If  the  substance 
is  insoluble,  decant  the  HC1,  wash  the  residue  twice  with  a  little  water  and 
decant  the  washings.  Then  try  the  action  of — 

Nitric  Acid.  Treat  the  residue  insoluble  in  water  and  HC1,  first,  with  boil- 
ing dilute  HN03  and,  secondly,  if  necessary,  with  concentrated  HN03  just  as 
in  the  extraction  with  HC1.  If  the  substance  is  soluble  in  HN03,  heat  the  solu- 
tion to  boiling  and  add  concentrated  HC1  as  long  as  brown  fumes  or  the  odor 
of  chlorin  is  evolved.  The  formation  of  a  precipitate  on  the  addition  of 
HC1  indicates  the  presence  of  Group  I,  Div.  A. 

By  the  interaction  of  HC1  with  HN03  (see  page  43),  the  latter  is  entirely 
removed  from  solution.  This  is  necessitated  by  the  fact  that  HN03  oxidizes  H2S 
to  free  sulfur  or  to  H2S04.  For  this  reason  its  presence  in  solution  would  inter- 
fere seriously  with  the  Basic  Analysis,  preventing  the  precipitation  of  sulfids  of 
the  metals  of  Group  I  and  possibly  causing  the  precipitation  of  sulfates  of 

239 


240  QUALITATIVE    ANALYSIS. 

Group  IV.  (For  these  same  reasons,  other  oxidizing  agents,  if  present,  interfere 
with  the  analysis.  This  method  of  boiling  with  concentrated  HC1  serves  to  re- 
move all  objectionable  oxidizing  agents,  however.)  "When  further  addition  of 
HC1  causes  no  further  liberation  of  brown  fumes,  evaporate  nearly  to  dry  ness  to 
remove  the  excess  of  acid.  Dilute  with  hot  water  and,  if  the  solution  becomes 
turbid  on  dilution,  add  sufficient  HC1  to  clear  it.  Filter,  if  necessary,  from  any 
Group  I,  Div.  A  precipitate.  Examine  this  precipitate  for  Ag,  Hg  and  Pb  as 
directed  under  Prec.  1,  page  250.  Treat  the  filtrate  as  directed  under  Filt.  1,  page 
252.  If  the  substance  does  not  dissolve  in  concentrated  HN03,  try  the  action  of — 

Aqua  Regia.  Add  concentrated  HC1  to  the  concentrated  HN03  and  warm. 
If  the  substance  dissolves,  remove  the  HN03  from  solution  and  proceed  as 
directed  in  the  preceding  paragraphs  devoted  to  the  extraction  with  HN03. 

Original  Solution.  The  solution  thus  prepared  for  systematic  analysis,  ob- 
tained with  or  without  the  use  of  acids,  is  known  as  the  original  solution. 

In  the  analysis  of  complex  mixtures,  the  preceding  treatment  may  result  in 
the  solution  of  one  or  more  components  of  the  mixture  by  each  of  the  different 
solvents  used.  When  partial  solutions  of  the  original  substance  are  thus  ob- 
tained, it  is  preferable  to  analyze  each  portion  separately;  for,  by  so  doing,  a 
better  idea  may  be  obtained  of  the  acid  which  is  united  to  any  given  base  present 
in  the  mixture.  If,  however,  it  is  found  that  the  separate  extracted  portions 
may  be  poured  together  without  producing  a  precipitate,  they  may  be  mixed 
and  the  whole  treated  as  the  original  solution  after  the  removal  of  nitric  acid 
as  previously  directed. 

During  the  treatment  with  HC1  in  making  the  original  solution,  indication 
is  frequently  obtained  of  the  presence  of  certain  acids.  When  ^carbonates,  sulfids, 
sulfites,  thiosulfates,  cyanids,  nitrites,  etc.,  are  treated  with  HC1,  gases  (C02, 
H2S,  S02,  HCN,  N204)  are  usually  liberated  which  may  be  recognized  by  their 
characteristic  odors  or  detected  by  the  regular  methods.  These  indications  should 
not  be  overlooked  and,  if  an  effervescence  is  noted,  the  gas  liberated  should  be 
determined  in  every  case. 

In  treating  with  HN03,  the  presence  or  absence  of  brown  fumes  denoting 
oxidization  should  also  be  noted. 

Metals  and  alloys  generally  should  be  dissolved  in  HN03,  for  treatment 
with  HC1  may  cause  the  loss  of  AsCl3,  H2S,  PH3,  etc.,  which  are  volatile. 
Phosphorus  and  sulfur,  present  in  an  alloy,  are  liberated  as  H2S  and  PH3  on 
treatment  with  an  acid  but  are  immediately  oxidized  by  HN03  to  non-volatile 
compounds: — free  sulfur,  H2S04,  H3P04,  etc.  If  gold,  platinum,  tin  or  anti- 
mony are  present,  the  alloy  will  not  dissolve  without  residue  in  HN03;  for  the 
first  two  metals  are  insoluble  and  the  last  two  form  meta-stannic  acid,  H2Sn03, 
and  antimony  pentoxid,  Sb205,  respectively,  compounds  quite  insoluble  in  nitric 
acid  as  well  as  in  water.  The  other  compounds  of  the  alloy  will  readily 
pass  into  the  solution  which  should  then  be  treated  as  previously  directed  for 
the  removal  of  HN03  and  then  tested  for  bases  in  the  regular  way.  None  of 
the  acid-forming  elements  need  be  tested  for,  except  sulfur,  phosphorus,  arsenic, 
silicon  and  carbon. 

The  residue  insoluble  in  HN03  should  be  washed  by  decantation  and  then 
digested  with  concentrated  HC1  for  10-15  minutes  at  the  boiling  temperature. 


THE    SYSTEMATIC  ANALYSIS.  241 

Add  an  equal  volume  of  water  and  again  heat  to  boiling.  This  treatment  should 
dissolve  most  of  the  tin  and  antimony  and  these  elements  may  be  detected  in 
the  decanted  solution  in  the  regular  way  by  passing  H2S  and  diluting  somewhat. 
The  residue  insoluble  in  HC1  should  be  treated  with  aqua  regia,  when  gold  and 
platinum  will  dissolve  and  may  be  detected  in  the  solution  in  the  usual  way. 
Should  any  residue  still  remain,  fuse  with  NaOH  and  sulfur  as  directed  later 
under  the  treatment  of  insoluble  substances  and  test  for  tin  and  antimony  in 
the  aqueous  extract. 

Insoluble  Substances.  The  substances  of  most  frequent  occurrence  which 
do  not  yield  to  the  foregoing  treatment  are :— ^  C,  AgCl,  AgBr,  AgL  AgCN", 
PbCrQ,,  PbS04,  Sn02,  &j&,  Sb204,  Fe^  Fe304,  Fe4(Fe  =  C6N8)3,  SrS04, 
BaS04,  A1203,  Cr  0,,  CaF,,  Si02  and  many  silicates.  These  substances  are 
usually  white  with  the  *  following  exceptions:  S  and  Agl  are  light  yellow; 
SnS2,  "mosaic  gold/'  is  deep  yellow;  PbCr6>4  and  Fe203  are  reddish;  Cr203  and 
sometimes  CaF2  are  greenish;  Fe4(Fe  =  C6N6)3  is  deep  blue,  Fe304  and  C  black. 

These  substances  require  special  treatment  varying  with  the  nature  of  the 
compound.  It  is  of  great  advantage  to  obtain  some  indication  as  to  the  com- 
position of  the  substance  before  treating  with  other  reagents.  Unless  sufficient 
indication  has  already  been  obtained,  examine  the  insoluble  residue  according 
to  the  methods  of  the  Preliminary  Examination  especially  as  directed  under  C. 

If  sulfur  or  carbon  is  detected,  it  must  be  removed  before  continuing 
the  analysis.  This  may  be  accomplished  most  conveniently  by  ignition.  Sulfur 
will  burn  to  S02  which  may  be  recognized  by  its  odor.  Carbonaceous  matter 
burns  to  C02,  the  black  substance  disappearing  entirely  or  leaving  only  a  slight 
white  ash. 

If  silver  is  indicated,  cover  the  substance  with  a  piece  of  metallic  zinc  and 
add  dilute  H2S04.  Continue  the  treatment  until  only  a  dark  residue  remains. 
The  zinc,  dissolving  in  the  H2S04,  throws  out  the  silver  in  metallic  form,  leav- 
ing the  acid  radicle  in  solution,  along  with  the  excess  of  zinc  sulfate,  where  it 
may  be  detected  in  the  usual  ways.  The  silver  may  be  obtained  in  solution  by 
dissolving  the  washed  residue  in  nitric  acid. 

If  lead  is  indicated,  fuse  the  substance  with  4-5  times  its  weight  of  a  mix- 
ture of  Na2C02  and  K2C08  on  a  piece  of  sheet  iron — platinum  may  be  ruined- 
cool,  extract  with  boiling  water  and  test  for  H2S04  and  H2Cr207.  The  lead 
may  be  obtained  in  solution  by  dissolving  the  washed  residue  in  acetic  or  nitric 
acid. 

If  tin  or  antimony  is  indicated,  fuse  the  substance  in  a  porcelain  crucible 
with  its  own  weight  of  NaOH  and  an  equal  amount  of  sulfur.  This  will  form 
soluble  thiostannates  or  thioantimonates  (see  pages  141  and  169)  and  in  the 
aqueous  extract,  treated  as  Filt.  5,  these  elements  may  be  detected  and  con- 
firmed in  the  regular  way. 

If  no  metal  other  than  iron  is  indicated  by  the  preliminary  test,  fuse  the 
substance  in  a  platinum  crucible  with  4-5  times  its  weight  of  the  Na2C03  — 
K2C03  mixture  and  add  a  few  crystals  of  KN03.  Cool,  boil  out  with  a  little 
water  and  test  the  aqueous  extract  for  the  acids  possibly  present;  also  for 
aluminum  and  chromium  which  are  present  (if  present  at  all),  the  former  as 
17 


242  QUALITATIVE    ANALYSIS. 

Na3A10s  and  the  latter  as  Na2O04.  Dissolve  the  insoluble  residue,  after  wash- 
ing, in  HC1  and  test  the  solution  for  the  bases. 

By  this  fusion,  cyanids,  either  simple  or  complex,  are  completely  oxidized, 
hence  this  acid  cannot  be  detected  in  the  resulting  solution.  All  cyanids  are 
decomposed  by  heating  with  dilute  H2S04  (1:2)  with  the  liberation  of  HCN 
which  may  be  collected  in  a  trap  containing  NaOH.  After  washing  out  the 
contents  of  the  trap  into  a  test-tube,  hydrocyanic  acid  may  be  detected  in  the 
usual  ways.  (See  page  294.) 

Some  silicates  and  ignited  oxids  are  not  decomposed  readily  by  fusion  with 
the  Na2C03 —  K2C03  mixture  but  yield  to  -the  action  of  fused  acid  potassium 
sulfate,  HKS04.  The  substance,  mixed  with  4-5  times  its  bulk  of  the  acid 
sulfate,  may  be  fused  in  a  test-tube  over  a  Bunsen  flame  for  about  one  half 
hour  at  a  temperature  just  high  enough  to  cause  a  slight  liberation  of  white 
fumes  of  sulfur  trioxid.  (See  page  188.) 

Alkalies  present  in  insoluble  silicates  can  not  be  determined  after  either  of 
these  methods  of  treatment  inasmuch  as  salts  of  the  alkalies  are  introduced 
during  the  process.  Resort  is  had  to  the  following  method.  The  finely  powdered 
silicate  is  mixed  with  an  equal  amount  of  NH4C1  and  5  times  as  much  CaC08 
and  heated  to  bright  redness  for  half  an  hour  or  longer  (see  page  135).  The 
aqueous  extract,  after  precipitation  of  calcium  by  (NH4)2C03  and  (NH4)2C204, 
is  evaporated  to  dryness  and  ignited  to  remove  ammonium  salts.  The  residue 
is  then  moistened  with  HC1  and  examined  for  alkalies  by  the  flame  tests. 

The  Prepared  Solution.  Before  beginning  the  Acid  Analysis,  bases  other 
than  the  alkalies  must  be  removed.  This  may  be  satisfactorily  accomplished  by 
adding  to  a  portion  of  the  original  solution  Na2C03  until  a  strong  alkaline  re- 
action is  produced.  A  precipitate  will  fall  which  is  boiled  in  the  solution  for 
fifteen  minutes  and  then  filtered  off.  By  this  treatment,  all  bases  except  arsenic, 
antimony  and  the  alkalies  are  removed  as  carbonates,  or  their  decomposition 
products,  and  the  acids  remain  in  solution  as  soluble  alkali  salts.  Any  insoluble 
salts,  other  than  carbonates,  which  at  first  might  be  precipitated  in  alkaline 
solutions,  are  subsequently  transposed  by  the  boiling  Na2C03  solution  (see  page 
68).  Thorough  boiling  should  not  be  neglected,  for,  otherwise,  some  acids  may 
be  filtered  off  and  lost  in  the  precipitate.  If  either  arsenic  or  antimony  is 
detected  in  the  Basic  Analysis,  before  adding  Na2C03  a  portion  of  the  original 
solution  is  made  strongly  acid  with  HC1  and  H,S  is  introduced  as  long  as  a 
precipitate  falls.  To  the  filtrate  from  this  solution,  Na2C03  is  added  and  the 
treatment  finished  as  directed.  The  solution  thus  made  ready  for  the  Acid 
Analysis  is  known  as  the  prepared  solution. 

It  is  advantageous  and  customary  to  make  the  Basic  Analysis  before  start- 
ing the  Acid  Analysis,  inasmuch  as  the  former  is  more  systematic  than  the 
latter  and  indications  of  certain  of  the  acids  present  are  frequently  obtained 
during  the  Basic  Analysis.  At  the  beginning  of  the  Acid  Analysis,  preliminary 
tests  are  made  indicating  the  presence  or  absence  of  certain  groups  of  acids. 
Further  than  this,  the  Acid  Analysis  consists,  generally,  in  making  individual 
tests  for  each  acid  that  has  not  been  excluded  by  the  preliminary  tests  from  the 
number  of  those  possibly  present.  A  consideration  of  the  solubility  of  the 
original  substance  will  exclude  certain  acids  at  once.  For  instance,  if  the 


PEELIMINARY    EXAMINATION. 

original  substance  is  soluble  in  water  and  barium  is  the  metal  present,  it  is 
unnecessary  to  make  tests  for  H2S04,  H3P04,  H2Si08  and  others,  for  the  barium 
salts  of  these  acids  are  not  soluble  in  water  as  a  glance  at  the  Table  of  Solu- 
bilities (pages  330  and  331)  will  show.  Similarly,  if  silver  has  been  found  in  a 
compound  forming  a  neutral  solution  with  water,  the  same  table  indicates  that 
the  only  acids  that  can  be  present  are  HF,  HN03,  HN02,  1I(C2H,02),  HC103, 
HC104,  HMn04  and  H2SiF6.  By  intelligent  use  of  this  Table  of  'Solubilities 
the  Acid  Analysis  may  usually  be  greatly  simplified  and  shortened. 

Liquids  submitted  for  qualitative  analysis  generally  consist  of  aqueous  solu- 
tions. After  noting  the  reaction  of  a  solution  to  litmus,  the  remainder  of  the 
preliminary  examination  may  usually  be  omitted.  If  the  solution  is  very  acid, 
remove  as  much  of  the  excess  of  acid  as  is  possible  by  evaporating  nearly  to 
dryness,  after  which  dilute  with  water  and  proceed  with  the  Basic  Analysis. 
If  the  solution  is  alkaline,  make  faintly  acid  w;th  HN03.  If  a  precipitate 
falls,  filter  and  consider  the  filtrate  as  the  original  solution  and  analyze  for 
bases  and  acids.  Treat  the  precipitate  (if  there  is  one)  according  to 
methods  for  insoluble  substances. 

Gases  can  be  satisfactorily  analyzed,  usually,  only  by  the  methods  of  Gas 
Analysis,  for  which  the  student  is  referred  to  the  standard  manuals  on  the 
subject. 

PRELIMINARY  EXAMINATION  OF  SOLIDS. 

For  these  tests  small  portions  of  the  original  substance  may  be  used.  If  the 
original  substance  be  in  solution,  its  reaction  to  litmus  should  be  noted.  Other- 
wise, the  preliminary  examination  may  usually  be  omitted  without  detriment. 
If  it  is  desirable  to  make  these  tests,  small  portions  of  the  residue  left  by 
evaporation  to  dryness  may  be  used.  Blackening  or  charring  on  evaporation  to 
dryness  is  an  indication  of  the  presence  of  organic  matter. 

Note  carefully  the  appearance  and  physical  properties  of  the  original  sub- 
stance and  draw  inferences  according  to  previous  experience. 

A.  Heat  in  a  Closed-tube  of  Hard  Glass.     (See  pages  53-55.) 
Observation.  Indication. 

1.  Change  of  color:  Presence  or  formation  of — • 

Yellow,  hot ;  straw-color,  cold ZnO. 

Yellow-red,  hot;  yellow,  cold;  fusible PbO. 

Yellow-brown,  hot;  dirty  white,  cold. ../...  Sn02. 

Orange,  hot;  light  yellow,  cold;  fusible Bi203. 

Dark  red,  hot;  yellow,  cold;  fusible PbCr04,  or  alkali  chromate. 

Yellow,  easily  fusible  to  brown  liquid S,  free  or  from  polysulfids. 

Red,  cold;  black,  hot;  globules  of  Hg HgO. 

Dull  red,  cold;  red-black,  hot Fe203. 

Becomes  black  on  heating — 

(a)  often    showing    blue    or    green    color 
previously    Co  or  Cu  salts. 

( b )  chars Organic  matter. 


244 


QUALITATIVE    ANALYSTS. 


Observation.  Indication. 

2.  Evolution  of  volatile  matter:  Presence  or  formation  of — 

Brown  color,  characteristic  odor . .  N204.  Nitrate  or  nitrite. 

Violet   color,   characteristic   odor,  lodin,  or  iodid  with  oxidizing 

black  sublimate I  agent. 

Brown-red     color,     characteristic  Bromid     with     an     oxidizing 

odor  Br.  agent. 

Pale    green    color,    characteristic  Chlorid     with     an     oxidizing 

odor   Cl.  agent. 

Colorless,  odor  of  burning  sulfur,  Sulfur,    sulfid,    sulfite,    thio- 

acid   reaction S0a.  cyanate,  sulfate  with  reduc- 

ing matter  or  other  sulfur 
compound. 

Colorless,  characteristic  odor H2S.  Sulfid,    or    sulfur    compound 

with  reducing  matter. 
Colorless,    making    "  lime-water "  Carbonate  or  organic  matter. 

turbid C02. 

Colorless,   characteristic   odor,   al-  Ammonium  salt. 

kaline  reaction NH3. 

Colorless,      characteristic      odor,  Cyanid  or  thiocyanate. 

burns  with  pink  flame (CN)2. 

Colorless,      odor      of      "burned  Organic  matter. 

sugar,"  inflammable. 
Colorless,  kindling  glowing  spark  0.  Chlorate,  nitrate,  peroxid  or 

unstable  oxid. 

Condenses  to  a  clear  liquid,  turn-  Moisture,  water  of  crystalliza- 

ing  dehydrated  CuS04  blue...H20.  tion  or  water  formed  dur- 

ing decomposition. 

—if  acid  in  reaction Volatile  acid  or  acid  salt. 

— if  alkaline  in  reaction Ammonium  salt. 


3.  Formation  of  a  sublimate: 
White  . 


White,  crystalline Sb203. 

White,  crystalline As203. 

White,  crystalline HgCl,. 

White,  yellow  when  hot (HgCl) 

Yellow As2S3. 

Yellow,  red  when  rubbed HgI2. 

Black As. 

Black,   violet   vapor I. 

Black,  red  when  rubbed Hg"S. 

Metallic  mirror  or  globules Hg. 


Presence  or  formation  of — 
Ammonium  salt  or  some  or- 
ganic substance. 
Antimonious  oxid. 
Arsenious  oxid. 
Mercuric  chlorid. 
Mercurous  chlorid. 
Arsenious  sulfid. 
Mercuric  iodid. 
Arsenic. 
lodin. 

Mercuric  sulfid. 
Mercury. 


PKELIMINAEY    EXAMINATION. 


245 


B.  Heat  with  Concentrated  Sulfuric  Acid  in  a 

Observation. 
Evolution  of  gas: 

Brown  color,  characteristic  odor.  ,N204. 
Violet   color,   characteristic   odor, 

black  deposit   I. 

Brown-red     color,     characteristic 

odor   Br. 

Pale    green    color,    characteristic 
odor Cl. 

Green-yellow  color,  characteristic 

odor,    explosive C102. 

Colorless,  odor  of  burning  sulfur  S02. 


Colorless,  odor  of  burning  sulfur 

accompanied  by  blackening. 
Colorless,       characteristic       odor 

darkens  lead-acetate  paper. . . .  H2S. 
Colorless,    making    "  lime-water " 

turbid C02. 

Colorless,  odor  of  bitter  almonds  HCN. 
Colorless,  characteristic  odor,  acid 

reaction  H(C2H302) 

Colorless,  characteristic  odor,  acid 

reaction  HC1. 

Colorless,  characteristic  odor,  acid 

reaction,  etches  glass HF. 

Colorless,  deposits  silicic  acid  on 

contact  with  water SiF4. 

Colorless,      burning     with     blue 

flame    .  .,CO. 


Test-tube.     (See  page  55.) 

Indication. 

Presence  or  formation  of — 
Nitrate  or  nitrite, 
lodin  or  iodid. 

Bromid. 

Hypochlorite,  chlorate,  or 
chlorid  with  an  oxidizing 
agent. 

Chlorate. 

Sulfite,  thiosulfate,  or  reduc- 
ing matter,  possibly  H2S, 
acting  on  the  sulfuric  acid. 

Organic  matter. 

Sulfid. 

Carbonate  or  oxalate. 

Cyanid. 
Acetate. 

Chlorid. 
Fluor  id. 

Silicofluorid,   or  fluorid  with 
silicate  or  silicic  anhydrid. 
Cyanid  or  oxalate. 


246 


QUALITATIVE    ANALYSIS. 


C.  Heat  on  Charcoal.     (See  page  55.) 

(a)   Without  other  reagents. 
Observation.  Indication. 

1.  The   substance    fuses    and   is   ab-  Salts  of  Li,  Na  or  K. 

sorbed  by  the  charcoal. 

2.  The  substance  is  infusible: 

Moistened  with  water  gives  alka-  BaO,  SrO,  CaO  or  MgO. 

line  reaction. 
Moistened  with  dilute  solution  of 

cobalt     nitrate     and     reheated 

turns    Blue.*  A1203. 

Green.          ZnO. 

3.  The  substance  deflagrates Nitrates  or  chlorates. 

4.  A  coating  forms  on  the  charcoal  :f 

White,  garlic  odor As203.  Arsenic. 

White,  near  the  residue Sb208.  Antimony. 

Yellow,  hot;  white,  cold ZnO.  Zinc. 

Yellow-red,  hot;   yellow,   cold...PbO.  Lead. 
Yellow-brown,   hot;    dirty   white, 

cold,  very  near  residue Sn02.  Tin. 

Orange,  hot ;  light  yellow,  cold . . .  Bi203.  Bismuth. 

Brown-orange,  hot;  same,  cold. . .  CdO.  Cadmium. 

(b)   With  Na2C03  and  powdered  charcoal. 
1.  Metallic  globule  is  formed  :$ 

Without    coating    on 

charcoal —  yellow,  Gold. 

red,  Copper, 

white,  malleable,    Silver. 

*  Some  silicates  give  a  pale  blue  color  on  similar  treatment. 

f  The  indications  given  by  the  oxids  here  formed  as  coatings  on  the  charcoal  may 
be  obtained  by  what  is  known  as  the  film  test.  A  small  portion  of  the  substance 
(dry  or  in  solution)  is  heated  on  a  few  fibres  of  asbestos  in  the  oxidizing  flame  of  a 
Bunsen  burner  so  placed  under  a  porcelain  dish  containing  cold  water  that  any  volatile 
matter  will  be  condensed  upon  it.  On  exposure  of  the  sometimes  invisible  film  to  H2S 
the  characteristic  color  of  the  corresponding  sulfid  will  be  developed. 

Film.  With  H2S.  Indication. 

White.  White.  Zinc. 

White.  Yellow.  Arsenic. 

White.  Orange.  Antimony. 

White.  Yellow  or  brown.  Tin. 

Yellow   (pale).  Brown.  Bismuth. 

Yellow.  Black.  Lead. 

Brown.  Yellow.  Cadmium. 

J  These  metallic  globules  may  be  very  satisfactorily  obtained  by  what  is  known 
as  the  match  test.  On  a  crucible  cover  fuse  a  piece  of  crystallized  sodium  carbonate 


PKELIMINAKY    EXAMINATION.  247 

With  coating  on  char- 
coal as  in  C  (a) 4 —  white,  malleable,    Tin  or  lead*  (see  pages  138 

and  142). 
white,  brittle,          Antimony  or  Bismuth  (see 

pages  165  and  170). 

2.  Magnetic  mass  is  dark,  brittle,          Iron,  cobalt  or  nickel, 

formed : 

D.  Heat  in  a  Borax-bead.     (See  page  56.) 
The  bead  is  colored: 
In  oxidizing  flame  —        In  reducing  flame — 

Blue.  Blue.  Cobalt. 

Blue.  Red,  opaque.  Copper. 

Green.  Green.  Chromium. 

Yellow.  Green.  Iron. 

Brown.  Gray,  opaque.  Nickel. 

Amethyst.  Nearly  colorless.  Mangaese. 

E.  Heat  on  Platinum  Wire  in  Bunsen  Flame.     (See  page  56.) 
The  flame  is  colored :  Yellow.  Sodium. 

Lavender.  Potassium. 

Orange.  Calcium. 

Crimson.  Strontium. 

Carmine.  Lithium. 

Green.  Copper  or  boric  acid. 

Green,  yellowish.  Barium. 

Blue,  brilliant.  Cupric  chlorid. 

Blue,  pale.  Arsenic. 

Na2CO3.10H2O,  and  rub  a  match  from  which  the  tip  has  been  removed  in  the  solution 
thus  formed  until  the  wood  is  saturated.  Char  the  match  in  a  Bunsen  burner,  dip  it  in 
the  finely  powdered  substance  and  heat  again  in  the  reducing  flame  when  the  metallic 
globule  will  appear. 

*  In  these  as  in  other  cases  where  it  is  necessary  to  determine  the  composition  of 
very  small  amounts  of  matter,  a  whole  series  of  reactions  may  be  performed  by  plac- 
ing the  substance  in  solution  and  making  separate  tests  on  single  drops  transferred 
by  means  of  a  glass  rod  to  a  porcelain  surface  or  filter  paper.  By  means  of  the  glasa 
rod  single  drops  of  reagents  (NaOH,  Na2CO3,  (NH4)2S,  NH4OH,  HC1,  H2SO4,  K2CrO4, 
KI,  KCN,  K4(Fer  C6N6),  K3(Fe  =  CGNG,  KSCN— some  of  these  or  others  as  suggested 
by  the  results  obtained)  may  be  added  and  the  character  of  typical  compounds  of  the 
metallic  component  noted.  The  formation  of  white  compounds  is  best  detected  by 
placing  the  drops  on  a  dark  surface.  This  together  with  the  film  test  and  match  test 
are  some  of  the  methods  of  microcliemical  analysis. 


248 


QUALITATIVE    ANALYSIS. 
THE  BASIC  ANALYSIS. 


Group  Separations.  —  The  following  scheme  indicates  the  manner  in  which 
the  base-forming  elements  are  divided  into  groups,  the  members  of  which  are 
individually  separated  and  identified  subsequently.  In  the  solution  to  be 
analyzed,  certain  of  the  following  list  of  kathions  are  present.  Any  changes 
caused  by  the  addition  of  reagents,  such  as  the  removal  of  certain  ions  by 
precipitation  in  the  form  of  insoluble  salts,  are  indicated  by  the  scheme. 


-H-         +++ 


I  I  I  I          ++        -H-H-         +-H- 


Ag,  CHg^2,  Pb,  Hg,  Bi,  Cu,  Cd,  As,  As,  +  Sb,  +Sb,    Sn,  Sn,    Al,   Cr,   Fe,  Fe, 
Co,  Ni,  Mn,  Zn,  Ba,  Sr,  Ca,  Mg,  Li,  Na,  K,  (NH4). 


HC1,  i.  e.,  Cl,  forms  Group  1,  Div.  A. 
AgCl,  (HgCl)2,  PbCl2. 


Pb,  Hg,  Bi,  jC^jCd,JLs,  As,  Sb^  Sb,  Sn, 
Sn",  Al,  Cr,  Fe,  Fe,  Co,  Ni,  Mn,  Zn,  Ba,  Sr, 
Ca,  Mg,  Li,  Na,  K,  (NH4). 


H,S,  i.  e.,  S,  forms  Group  I,  Divs.  B  and  C. 
PbS,  HgS,  BIS,,  CuS,  CdS,  AsaSs, 
As,S5,  SbaS3,  Sb^S.,  SnS,  SnS2. 
(NH4)2S.xS  gives  a— 
Res.  —  Group  I,  Div.  B. 

PbS,  HgS,  Bi.8,,  CuS,  CdS. 
Filt.—  Group  I,  Div.  C. 
(NH4)3AsS4,  (NH4)3SbS4 


HC1,  i.  e.,  H,  forms  Group  I,  Div.  C. 
As2S5,  Sb8SB, 


Al,    Cr,   Fe,  Co,  Ni,  Mn,  Zn,  Ba,  Sr, 
Ca,  Mg,  Li,  Na,  K,  (NH4). 

Boil  with  HN03. 

Al,  Cr,  Fc,  Co,  Ni,  Mn,  Zn,  Ba,  Sr, 
Ca,  Mg,  Li,  Na,  K,  (NHJ. 


NH4(C2H302)  and  boiling  forms  Group  II. 
A1(OH)2(C2H3(X),  Cr(OH)2(C2H302), 
Fe(OH)2(C2HA). 


(NH4)2S,  ».  e.,  S,  forms  Group  III. 
CoS,  NiS,  MnS,  ZnS. 


Co,  Ni,  Mn,  Zn,  Ba,  Sr,  Ca,  Mg, 
L+i,  Na,  K,  (NH4). 


Ba,  Sr,  Ca,  Mg,  Li,  Na,  K, 
(NHJ. 


(NH4)2C03,  t.  e.,  COS,  forms  Group  IV. 
BaC03,  SrC03,  CaC03. 


Mg,  Li,  Na,  K, 
(NH4). 


HNa2P04,  f.  e.,  P04,  precipitates  (one  metal  only  of)  Group  V. 
Mg(NH4)P04. 


Li,  Na,  K, 
(NH4). 


THE   BASIC    ANALYSIS.  249 

SPECIAL  TESTS. 

Certain  base-forming  substances,  such  as  ammonium  and  iron,  are  introduced 
as  reagents  into  the  solution  to  be  investigated  during  the  course  of  the  analysis. 
Others,  originally  present  in  one  condition,  are  likely  to  be  changed,  i.  e.,  oxidized 
or  reduced,  during  the  process.  It  is  necessary,  therefore,  to  test  for  the  presence 
and  condition  of  certain  substances  before  beginning  the  systematic  Basic  Analysis. 


Test  for  Ammonium,  (NH4). 

Heat  a  small  portion  of  the  original  dry  substance  with  dry  Ca(OH)2,  or  add 
to  a  small  portion  of  the  original  solution  NaOH  or  KOH  until  the  solution  is 
strongly  alkaline  and  boil.  The  liberation  of  pungent  fumes  of  ammonia,  NH3, 
which  turn  moist  red  litmus-paper  blue,  indicates  (NH4).  (See  page  76.) 


Test  for  the  Presence  and  Condition  of  Iron,  Fe. 

(1)  Acidify  a  small  portion  of  the  original  solution  with  HC1  and  add  a  drop 


Dark  blue  prec.,  Fe4(Fe  =  C6N6)3,  indicates  Fe,  present  in  the  trivalent  or 
ferric  condition. 

(2)  Acidify  a  small  portion  of  the  original  solution  with  HC1  and  add  a  drop 
of  K3(Fe==C6NJ  solution,  freshly  prepared  by  dissolving  a  small  crystal  in 
water.  Dark  blue  prec.,  Fe3(Fe=C6N0)2,  indicates  Fe,  present  in  the  bivalent  or 
ferrous  condition.  (See  page  227.) 


Test  to  Determine  the  Condition  of  Tin,  Sn. 

Acidify  a  small  portion  of  the  original  solution  with  HC1,  filter,  if  necessary 
to  obtain  a  clear  solution  and  add  a  few  drops  of  HgCl2.  (See  page  110.) 

White  prec.,  (HgCl)^  which  may  turn  dark  on  standing,  Hg,  indicates  Sn, 
present  in  the  bivalent  or  stannous  condition. 

If  tin  is  found  during  the  Basic  Analysis  and  no  prec.  is  obtained  here,  it  must 
be  present  in  the  tetravalent  or  stannic  condition. 


250 


QUALITATIVE    ANALYSIS. 


SEPARATION  AND  ANALYSIS   OF  GROUP  I,  DIV.  A. 

(If  HC1  was  used  in  making  the  solution  (see  page  239),  proceed  at  once  to 
the  Separation  of  Group  I,  Divs.  B  and  C,  page  252.) 

To  the  original  solution  add  one-tenth  its  volume  of  cone.  HC1.  Shake  thor- 
oughly, filter,  and  wash  the  pree.  with  cold  water. 


Prec.  1.  Group  I,  Div.  A. 
AgCl,  (HgCl)2,  PbCl2. 


FILT.  1.  Group  I,  Divs. 
B  and  C  and  Groups  II- 
V.  See  page  252. 


Treat  on  the  filter  paper  thoroughly  with  boiling  water  and  allow  the  washings 
to  filter  through  into  a  test-tube  or  beaker. 


RES.  2.    AgCl,  (HgCl)2, 

Wash  thoroughly  with  boiling  water  until  the 
washings  no  longer  give  a  test  for  lead. 

Reject  the  washings. 

Moisten  thoroughly  with  2-3  ccm.  NH4OH,  avoid- 
ing any  excess  of  the  reagent,  and  wash  with  a  little 
water. 

RES.  3.     Hg2NH2Cl. 

Black  coloration  indi- 
cates Hg. 

Treat  on  the  filter 
paper  with  a  little  warm 
aqua  regia  and  wash  into 
a   test-tube.      (Filter   if 
necessary'  to  remove  any 
AgCl  or  filter  fibre.)     To 
the  clear  filt.  add  SnCl,. 
White  prec.,    (HgCl)2,~ 
or  dark,  Hg,  confirms 
Hg. 


FILT.  3. 

Acidify  with  HN03. 
White  prec.,  AgCl,  dark- 
ening in  sunlight,  indi- 
cates Ag.* 


FILT.  2. 

Divide  into  three  por- 
tions. 

A.  Add  2  or  3  drops 
K2Cr04.     Yellow  prec., 
PbCr04,  soluble  iri 
NaOH,  indicates  Pb. 

B.  Add  H2S04  and 
allow  to  stand  for  a 
few  minutes  if  necessary. 
White  prec.,  PbS04, 
soluble  in  NaOH,  indi- 
cates Pb. 

C.  Add  Kl.f    Yellow 
spangles,  PbI2,  soluble  in 
hot  water,  recrystallizing 
on  cooling,  indicate  Pb. 


*  If  no  indication  of  the  presence  of  Ag  is  obtained  here  and  a  blackening  indi- 
cating mercury  was  noted  in  the  proper  place,  it  is  safest  to  test  the  Hg  prec.  for  Ag 
also.  After  treatment  with  aqua  regia,  any  Ag  present  will  be  in  the  form  of  the 
chlorid  which  may  be  filtered  off  (see  treatment  of  Res.  3)  and  its  solubility  tested  in 
NH4OH,  KCN  or  Na2S2O3. 

t  A  few  drops  only,  as  an  excess,  will  form  the  soluble  double  salt,  K2PbI4  (see 
page  145). 


THE    BASIC    ANALYSIS.  251 

DISCUSSION  OF  THE  SEPARATION  AND  ANALYSIS  OF 
GROUP  I,  DIV.  A. 

The  separation  of  Group  I,  Div.  A  from  the  salts  of  other  metals  which  may 
be  present  in  solution  depends  on  the  relative  insolubility  of  the'  chlorids  of  these 
elements  and  the  solubility  of  other  chlorids.  The  separation  is  not  exact,  how- 
ever, since  PbCl2  is  very  appreciably  soluble  and  will  not  be  precipitated  at  all 
from  solutions  containing  one  per  cent,  or  less  of  the  Pb  salt.  For  this  reason  Pb, 
if  present,  will  always  be  found  in  Group  I,  Div.  B  but  may  not  be  detected  in  Div. 
A.  On  the  other  hand,  the  oxychlorids  of  Sb  and  Bi  (SbOCl  and  BiOCl)  are 
quite  insoluble  and  may  be  precipitated  by  a  small  amount  of  HC1.  These  salts 
will  dissolve  in  an  excess  of  acid,  however,  hence  the  reason  for  adding  cone.  HC1 
to  the  extent  of  one-tenth  the  volume  of  the  solution. 

Mercury  as  well  as  Pb  may  be  found  in  both  Divs.  A  and  B  of  Group  I.  The 
chlorid  formed  by  the  mercuric  ion,  Hg,  is  soluble,  hence  salts  of  mercury  in  the 
divalent  condition  are  separated  from  the  univalent  mercurous  compounds,  for 
the  chlorid  formed  by  the  mercurous  ion,  (Hg)2,  is  highly  insoluble. 

The  -solubility  of  PbCl2  in  boiling  water  offers  an  easy  means  of  separation 
from  Ag  and  Hg  in  the  lower  condition,  the  chlorids  of  these  metals  being  prac- 
tically insoluble.  Under  certain  conditions,  the  prec.  of  (HgCl)2  is  so  finely 
divided  as  to  wash  through  the  filter  paper  to  some  extent  and  give  a  cloudy  filt. 
If  Pb  is  absent,  or  present  in  small  amount  only,  a  reddish  coloration  due  to 
Hg2Cr04  may  appear  on  adding  the  Cr04  ion.  Since  the  Hg  salt  is  insoluble  in 
NaOH  or  KOH,  it  may  be  distinguished  from  the  Pb  salt  which  dissolves  readily 
with  the  formation  of  an  alkaline  plumbite  solution  containing  the  ion  Pb02. 
(See  page  144.) 

The  formation  of  complex  ions  on  treatment  with  NH4OH  is  the  basis  of  the 
separation  or  Hg  and  Ag,  since  the  compounds  formed  by  the  ion  Ag(NH3)2  are 
soluble  while  compounds  formed  by  (Hg),NH,  are  insoluble.  Salts  of  this  latter 
type  decompose  immediately  into  metallic  Hg  and  compounds  containing  the 

radicle    (  — Hg — N<  ii  j .    Since  the  Ag  ion,  and  more  readily  the  Ag(NH3)2  ion, 

is  reduced  to  the  metal  by  Hg,  in  the  presence  of  a  large  amount  of  the  Hg  salt 
a  small  amount  of  Ag  salt  may  be  entirely  decomposed  in  this  way  and  remain 
as  metallic  Ag  in  the  black  residue.  A  large  excess  of  NH4OH  is  to  be  avoided, 
therefore.  The  mercury  compound  may  be  decomposed  and  the  metal  driven  off 
by  igniting  in  porcelain  under  a  hood,  after  which  the  Ag  may  be  dissolved  in 
HN03  and  tested  for  in  the  usual  way.  Or,  the  black  residue  may  be  dissolved  in 
aqua  regia,  by  the  action  of  which  all  Hg,  both  free  and  combined,  is  oxidized  to 
HgCl2.  This  may  be  separated  from  the  insoluble  AgCl  by  filtration  (see  foot-note 
on  opposite  page).  Stannous  chlorid,  SnCl2,  being  a  strong  reducing  agent,  re- 
duces HgCl2  first  to  white  (HgCl)2  and  finally  to  Hg  which  in  finely  divided 
condition  is  black. 

Should  Filt.  3  be  milky,  it  is  probably  due  to  the  fact  that  all  PbCl2  was  not 
washed  out  of  Res.  2.  With  NH4OH,  PbCl2  forms  a  basic  chlorid  which  is  very 
apt  to  run  through  the  filter.  It  will  dissolve  in  HN03,  however,  hence  will  not 
interfere  with  the  test  for  Ag. 


252 


QUALITATIVE    ANALYSIS. 


SEPARATION  OF  GROUP  I,  DIVS.  B  AND  C. 

Dilute  Filt.  1  with  an  equal  volume  of  water,  warm  and  pass  H2S  until  precipi- 
tation ceases.  Filter  and  test  a  portion  of  the  nit.  by  diluting  with  three  volumes 
of  water  and  again  passing  H2S.  Should  precipitation  occur  again,  dilute  the 
whole  nit.  and  pass  H2S  again  until  precipitation  is  complete.  Filter  and  wash 
with  hot  water. 


PREC.  4.     Group  I,  Divs.  B  and  C. 
HgS,  PbS,  ]*03,  CuS,  6*6,  Sb2S3,  Sb2£ 
SnS,  SnS2  (AfeA^Jtoft,). 

Puncture    the    filter    paper    and   wash   the    prec. 


FILT.  4.    Groups  II-V. 
Boil  out  H2S*  imme- 
diately and  treat  as  di- 
rected on  page  260. 


into  a  test-tube  or  small  beaker  with  5-10  ccm.  of  yellow  ammonium  sulfid, 
(NH4)2S.xS.    Warm,  filter  and  wash. 


RES.  5.     Group  I. 
Div.  B. 

Analyze  as  directed  on 
page  254. 


FILT.  5.     Group  I,  Div.  C. 
Acidify  with  HC1  and  filter. 

PREC.  6.     Group  I,  Div.  C.  FILT.  6.    Reject. 

As,S5,  Sb2S5,  SnS2,  S, 
(Au2S3,  PtSJ. 

Analyze  as  directed  on  page 
256. 


DISCUSSION  OF  THE  SEPARATION  OF  GROUP  I,  DIVS.  B  AND  C. 

The  separation  of  Group  I  from  the  other  groups  of  metals  which  may  be 
present  in  solution  depends  on  the  fact  that  the  sulfids  of  Ag,  Hg,  Pb,  Bi,  Cu,  Cd, 
As,  Sb  and  Sn  are  relatively  insoluble  in  water  and  dilute  acid.  The  sulfids  of 
Group  II  and  of  Mg  are  hydrolyzed  completely  in  aqueous  solution  with  the 
precipitation  of  hydroxids  insoluble  in  water  but  soluble  in  dilute  acid.  The 
sulfids  of.  Group  III  are  either  soluble  in  dilute  acid  or  their  precipitation  is 
prevented  by  it.  The  sulfids  of  Group  IV  and  V  are  soluble.  Hence,  in  the 
presence  of  dilute  acid,  only  the  sulfids  of  Group  I  are  precipitated. 

The  metals  of  Group  I,  Div.  A  are  not  ordinarily  precipitated  as  sulfids,  how- 
ever. Because  of  the  relative  insolubility  of  the  chlorids  of  Ag,  Hg(ous),  and 
Pb,  Div.  A  is  separated  after  the  addition  of  HC1  and  before  the  addition  of  the 
sulfid  ion.  Consequently  Divs.  B  and  C  only  are  ordinarily  precipitated  as  sulfids. 

Since  H'S  is  a  reducing  agent,  if  oxidizing  agents  be  present  in  solution,  an 
interaction  will  take  place  resulting  in  the  oxidation  of  H2S  to  H20  and  S,  or  to 
H2S04.  Under  these  circumstances  sulfids  are  precipitated  only  after  all  oxidizing 
agents  have  been  reduced.  If  present  in  considerable  amounts,  the  products  of 
this  reaction,  mostly  S,  may  lead  to  difficulties  and  cause  confusion  in  the  analysis 
of  the  precipitate. 

Hence  it  is  better  to  remove  any  oxidizing  agent  before  beginning  the  analysis 
as  directed  on  page  239.  If  present  in  small  amounts  only,  they  will  not  interfere 
but  may  give  an  indication  of  some  of  the  elements  present.  A  pink  or  purple 

*  The  point  at  which  H2S  has  been  entirely  removed  is  indicated  by  the  fact 
that  a  filter  paper,  moistened  with  Pb(C2H3O2)2  and  held  in  the  vapor  from  the  solution, 
is  no  longer  blackened. 


THE    BASIC    ANALYSIS.  253 

solution  of  a  permanganate  is  bleached  as  the  H2S  enters.  An  orange  chromate 
solution  is  turned  green  and  S  is  precipitated.  However,  if  Pb  is  present  in  small 
amounts  only,  the  formation  of  PbS04  from  the  oxidation  of  H2S  may  cause  some 
difficulty.  In  this  case,  Res.  7  should  be  tested  for  Pb  as  well  as  Hg. 

The  sulfids  of  Div.  C,  especially  the  As  sulfids,  may  be  most  satisfactorily  pre- 
cipitated in  the  presence  of  considerable  acid.  This  prevents  the  formation  of 
soluble  salts  of  the  thioacids  and  of  colloidal  precipitates  at  the  same  time. 
Hence  it  is  better  not  to  dilute  Filt.  1  too  extensively  before  passing  H2S. 
On  the  other  hand,  PbS  is  soluble  in  fairly  strong  HC1  and  CdS  in  a  concentra- 
tion greater  than  4  per  cent.  The  precipitation  is  made  in  a  hot  solution  to  hasten 
the  reactions  and  prevent  colloidal  precipitates  which  sulfids  as  a  class  show  strong 
tendency  to  form.  The  final  dilution  and  testing  of  the  filtrate  as  directed  is 
essential  to  insure  the  precipitation  of  CdS  which,  if  not  removed  here,  will 
come  down  with  Group  III.  Thorough  washing  of  Prec.  4  is  necessary  to  remove 
members  of  other  groups,  as  otherwise  Al  or  Fe  may  interfere  with  the  test  for 
Bi  as  indicated  in  foot-note  on  page  254. 

It  is  absolutely  necessary  that  Filt.  4  be  boiled  free  from  H2S  immediately, 
for  otherwise,  the  oxygen  of  the  air  will  convert  it  into  H2S04  and  cause  the 
precipitation  of  sulfates  of  Group  IV  (Res.  21). 

The  separation  of  Divs.  B  and  C  from  each  other  depends  on  the  solubility  of 
the  sulfids  of  Div.  C  in  alkaline  polysulfids  and  the  relative  insolubility  of  the 
sulfids  of  Div.  B.  Ammonium  polysulfid,  (NH4)2S.  xS,  gives  the  sharpest 
separation  but  even  this  is  not  exact,  for  if  a  large  excess  of  the  reagent  is  used, 
CuS  may  be  dissolved  in  part  and  pass  into  the  filtrate  to  interfere  with  the  test 
for  Sb.  (See  foot-note,  page  256.) 

Colorless  alkaline  sulfids  of  the  type,  (NH4)2S  dissolve  As2S5,  Sb2SB  and  SnS2 
to  form  soluble  thioarsenates,  thioantimonates  and  thiostannates  (see  pages  164, 
169,  141)  and  As2S3  and  Sb2S3  will  dissolve  in  the  same  reagent,  forming  soluble 
thioarsenites  and  thioantimonites.  Stannous  sulfid,  SnS,  does  not  form  a  corre- 
sponding thiostannite,  however,  and  will  not  dissolve  in  colorless  (NH4)2S.  For 
this  reason  it  is  necessary  to  use  the  yellow  polysulfid,  (NH4)2S.  x  S.  The  addi- 
tional S  in  compounds  of  this  type  oxidizes  SnS  to  SnS2,  which  then  dissolves 
to  form  the  thiostannate.  At  the  same  time,  the  sulfids  of  any  trivalent  As  or  Sb 
present  are  oxidized  to  the  pentavalent  condition,  thus  furnishing  an  example  of 
the  oxidizing  action  of  S. 

Whether  any  of  these  thioacids  are  capable  of  existence  in  the  free  state  is 
very  questionable.  At  any  rate  they  decompose  immediately  into  H2S  and  the 
metallic  sulfid  (for  reactions  see  page  165).  Consequently  on  acidifying  a 
solution  of  their  salts,  As2S5,  Sb2S5  and  SnS,  are  precipitated.  Even  if  present 
originally  in  the  lower  condition,  after  oxidation  by  the  polysulfid,  only  the  higher 
sulfids  are  precipitated  on  acidifying. 

Inasmuch  as  all  polysulfids  separate  free  S  when  treated  with  acids  (see 
page  184),  some  precipitate  is  always  obtained  in  this  place  whether  Div.  C  be 
present  or  not.  The  sulfids  of  these  elements  are  deep  yellow  or  orange  in  color 
and  are  flocculent  in  character,  while  the  precipitated  S  is  usually  straw-colored, 
very  finely  divided  and  does  not  tend  to  gather  together  to  any  great  extent,  giving 
rather  a  milky  solution  which  passes  through  a  filter. 


254 


QUALITATIVE    ANALYSIS.. 


ANALYSIS  OF  GROUP  I,  DIV.  B. 


RES.  5.  From  page  252.  Group  I,  Div.  B.  HgS,  PbS,  Bi2S3,  CuS,  CdS. 
Puncture  the  filter  paper  and  wash  the  res.  into  a  test-tube  with  5-10  ccm. 
water.  Add  one-half  its  volume  of  cone.  HN03  and  boil  for  two  or  three  minutes. 
Dilute  with  an  equal  volume  of  water  and  filter. 


RES.  7. 

FILT.  7.     Add  a  few  drops  of  cone.  H2S04  and  evaporate  until 

Black  HgS 

dense  white  fumes  of  S03  appear.    Cool,  dilute,  and,  if  the  solu- 

or white 

tion  is  not  distinctly  acid,  add  2-3  ccm.  H2S04  and  filter. 

HgS.Hg(N03)2 

or  possibly 

RES.  8. 

FILT  8.    Add  NH4OH  until  solution  is  alka- 

PbS04. 

White 

line  and  filter. 

Dissolve  in 

PbS04. 

a  small 

Confirm 

PREC.  9. 

FILT.  9.     Blue  solution  indi- 

amount of 

by  dissolv- 

White* 

cates  Cu.     If  Cu  is  indicated, 

aqua  regia, 

in  a  few 

Bi(OH)3. 

test  for  Cd  as  in  B  (1). 

dilute,  boil 
to  remove 
"chlorin  and 
filter.  To  the 
filt.  add 

drops  of 
NaOH.   Add 
K2Cr04  and 

acidify 
faintly  with 

Add  a  few 
drops  of  dil. 
HC1  to  the 
prec.  on  the 
paper  and 

If  no   color  appears,   divide 
filt.     Test  one  portion  for  Cu 
as  in  A,  and  another  for  Cd  as  in 
B  (1)  or  B  (2)  according  as  Cu 
is  or  is  not  indicated  in  A. 

SnCl2. 

H(C2H302) 

allow  to  filter 

A.  Acid- 

B   (1).     If  Cu 

White  prec., 

Yellow 

through  into  a 

ify  with 

is  indicated,   add 

(HgCl)2  or 
dark,  Hg, 

prec., 
PbCr04, 

beaker  of 
water. 

H(C2H302) 
and  add 

KCN     somewhat 
in   excess   of  the 

indicates 
the  presence 
of  Hg. 

indicates 
Pb. 

White  prec. 
BiOCl  indi- 
cates Bi. 
Or,  wash 

K4(Fe=C6N6). 
Mahogany- 
colored  prec., 

amount  necessary 
to  cause  the  blue 
color     to     disap- 
pear.    Pass  H,S. 

p,  jc.  on  the 
filter  and  add 

indicates  Cu. 

Yellow    prec.,:}: 
CdS,  indicates  Cd. 

a  few  drops  of 

(2)  If  Cu  is  not 
indicated,  acidify 

-• 

Na2Sn02.f 

slightly  with  HC1 

Dark  color- 

and pass  H2S. 

ation,  BiO, 

Yellow       prec., 

indicates  Bi. 

CdS,  indicates  Cd. 

*  Sometimes  slightly  colored  brown  by  traces  of  Fe(OH)3  if  Prec.  4  is  not  thor- 
oughly washed.  For  the  same  reason  a  slight  white  prec.  which  will  not  give  confirma- 
tory tests  for  Bi  may  be  A1(OH)3. 

t  Sodium  stannite  (NaL,SnO2)  solutions  must  be  freshly  prepared  each  time  just 
previously  to  using.  A  few  drops  of  SnCl2  solution  are  treated  with  NaOH  or  KOH 
with  vigorous  shaking  until  the  prec.  first  formed  just  redissolves,  thus  avoiding  any 
excess  of  alkali.  (See  page  139.) 

$  If  a  dark  prec.  falls,  it  is  due  to  traces  of  Pb  or  Hg  and  possibly  Bi  but  is  not 
due  to  Cu  if  sufficient  KCN  has  been  added.  Filter  off  the  prec.,  puncture  the  paper  and 
wash  into  a  test-tube  with  10  ccm.  water.  Add  2  ccm.  cone.  H2SO4,  boil  and  filter. 
Dilute  the  filt.  with  an  equal  volume  of  water  and  pass  H2S,  neutralizing  part  of  the 
acid  present  with  NH4OH  if  necessary  to  bring  down  a  prec.  Yellow  prec.,  CdS, 
indicates  Cd. 


THE    BASIC    ANALYSIS.  255 

DISCUSSION  OF  THE  ANALYSIS  OF  GROUP  I,  DIV.  B. 

Although  the  sulfids  of  Div.  B  are  not  caused  to  pass  into  solution  by  the  pres- 
ence of  the  H  ion,  with  the  exception  of  HgS  they  do  dissolve  in  dilute  HN03  be- 
cause of  its  oxidizing  action.  The  separation  of  Hg  from  the  other  members  of 
this  Div.  depends  upon  this  fact,  yet  it  may  sometimes  be  not  altogether  exact. 
Some  of  the  complex  Hg  salts,  at  first  precipitated  by  H2S  from  acid  solutions 
(see  page  108),  are  soluble  in  HN03  and,  if  not  entirely  changed  into  the  simple 
HgS,  will  cause  Hg  to  pass  into  Filt.  7  to  form  a  black  precipitate  later  and  inter- 
fere with  the  test  for  Cd.  Treatment  with  (NH4)2S.  x  S  insures  complete  trans- 
formation of  these  intermediate  Hg  compounds  into  HgS.  If,  for  any  reason,  the 
regular  separation  of  Divs.  B  and  C  by  means  of  this  reagent  is  omitted,  it  is 
advisable  to  moisten  the  precipitated  sulfids  with  a  little  diluted  (NH4)2S  x  S. 

The  oxidation  of  these  sulfids  by  means  of  HN03  results,  for  the  most  part,  in 
free  S  but  some  H2S04  may  be  formed,  especially  if  the  acid  be  stronger  than  5N 
(i.  e.  1  ccm.  cone.  HN03  to  2  ccm.  H20),  or  if  the  boiling  be  prolonged.  Either 
cause  may  precipitate  some  PbS04  which  may  remain  with  the  HgS  residue. 

The  separation  of  Pb  depends  on  the  fact  that  its  sulfate  is  the  only  one 
of  this  Div.  relatively  insoluble.  Inasmuch  as  PbS04  dissolves  very  appreciably 
in  HNO^  it  is  necessary  to  remove  the  latter  to  get  even  an  approximate  separa- 
tion. The  B.  P.  of  cone.  HN03  is  120° ;  the  fuming  point  of  H2S04  about  250°. 
It  is  fairly  certain,  therefore,  that  all  HN03  will  have  been  removed  when  the 
solution  has  been  heated  until  dense  fumes  of  S03  appear.  Even  after  this 
treatment,  however,  traces  of  Pb  may  be  found  in  Filt.  8  which  may  darken 
the  precipitate  of  CdS  later  in  the  analysis. 

The  separation  of  Bi  from  Cu  and  Cd  depends  upon  the  fact  that  in  the 
presence  of  NH4OH,  Bi(OH)3  is  precipitated  whereas  Cu  and  Cd  form  complex 
ammonia  ions,  the  compounds  of  which  are  soluble.  The  insoluble  Bi(OH)3  is 
transformed  by  HC1  into  soluble  BiCl3  which,  on  being  largely  diluted,  partially 
hydrolyzes  and  precipitates  BiOCl  as  a  faint  white  cloud.  More  than  a  very 
few  drops  of  HC1  must  be  avoided,  since  the  presence  of  even  a  small  con- 
centration of  H  ions  prevents  the  formation  of  this  precipitate.  For  the  dis- 
cussion of  the  confirmatory  test  for  Bi,  see  page  171. 

The  separation  of  Cu  from  Cd  depends  upon  the 'relative  stability  of  the 
complex  cyanid  ions,  the  formation  of  which  is  indicated  by  the  bleaching  out  of 
the  blue_color  due  to  the  ion  Cu(NH8)4.  The  (CuCjSTJ  ion  is  very  stable  while 
the  CdCjN,  ion  is  decomposed  by  H,S  with  the  precipitation  of  CdS.  For 
further  discussion  of  this  separation,  see  pages  81  and  84. 

Very  minute  traces  of  black  sulfids  other  than  Cu  may  darken  the  CdS 
precipitate  so  as  to  leave  the  test  indefinite.  Since,  however,  CdS  is  the  only 
sulfid  of  this  Div.  soluble  in  boiling  6N  H2S04  (i.  e.,  1  ccm.  cone.  H2S04  to 
5  ccm.  H,0),  a  thoroughly  satisfactory  separation  may  be  effected  by  this 
reagent.  From  the  filtrate  pure  yellow  CdS  may  be  precipitated  by  passing 
H2S  and  neutralizing  part  of  the  acid  with  a  few  drops  of  NH4OH. 

Copper  present  in  amounts  too  small  to  give  a  satisfactory  blue  color^with 
NH4OH  may  frequently  be  detected  by  means  of  the  ferrocyanid  ion,  (Fe  =  C6N6), 
which,  with  the  Cu  ion,  gives  a  reddish  precipitate  of  great  intensity  of  color. 
If  Cd  be  present,  white  Cd,(Fe  =  C6N6)  will  be  precipitated  which  may  dilute 
the  color  given  by  traces  only  of  Cu  so  that  the  precipitate  appears  flesh-colored. 


256 


QUALITATIVE    ANALYSIS. 


ANALYSIS  OF  GROUP  I,  DIV.  C. 

PEEC.  6.    From  page  252.    GroupI,Div.C.    As2S5,  Sb2S5,  SnS2(Au2S3,  PtS2). 

Puncture  the  filter  and  wash  the  prec.  into  a  test-tube  with  about  10  ccm.  cone.  HC1, 

dilute  with  half  its  volume  of  water,  boil  until  H2S  is  entirely  removed  and  filter. 

RES.  10.    As2S5  (Au2S3,  PtS2).  FILT.  10.     Divide  into  two  portions. 

A.  If  the  res.  is  yellow,  i.  e., 
contains  no  Au  or  Pt.     Wash 
the  prec.  on  the  filter  until  the 
washings  give  no  test  for  the  Cl 
ion.*     Puncture  the  filter  and 
wash  the  prec.  into  a  test-tube 
with  5-10  ccm.  HN03  and  boil 
thoroughly,  evaporating  nearly 
to  dryness  at  least  twice,  replac- 
ing   the    liquid    evaporated   by 
HN03.     Boil  out  the  res.  from 
the  last  evaporation  with  5-10 
ccm.  water.     Add  AgN03  and 
shake     thoroughly.      Carefully 
add  to  the  test-tube  2-3   cem. 
NH4OH  without  shaking. 

Formation  of  a  chocolate 
prec.,  Ag3As04,  in  the  neutral 
ring  between  the  acid  and  alka- 
line solutions  indicates  As. 

B.  //  Ees.  10  is  dark,  i.  e., 
probably    contains    Au    or    Pt, 
proceed  as  directed  on  page  258. 

*  Acidify  5  ccm.  of  the  washings  with  HNO3  and  add  a  drop  of  AgNO3.  White 
prec.  or  cloudiness,  AgCl,  indicates  the  CT  ion. 

t  The  most  satisfactory  results  may  be  obtained  by  keeping  the  contact  of  the 
metals  outside  of  the  solution  to  be  tested.  This  may  be  readily  accomplished  by 
bending  a  strip  of  Zn  into  the  form  shown  in  Fig.  1.  By  placing  the  solution  to  be 


A.  Make    a    little 
galvanic  couple  by 
bringing  into  contact 
a  piece  of  platinum 
and  another  of  zinc.f 
Immerse  the  metals 
in  one  portion  of 
Filt.  10  to  which  an 
equal  volume  of  cone. 
HC1  has  been  added. 

Black  deposit;$ 
metallic  Sb,  on  the 
platinum  electrode 
indicates  Sb. 


B.    To  the   second 
portion   add  3   or   4 
iron  nails  (not 
tinned)  and  allow  the 
action  to  go  on  for  10 
to  15  minutes  or  until 
the  solution  is  color- 
less, keeping  up  a 
steady  evolution  of 
hydrogen  by  the  addi- 
tion of  cone.  HC1  if 
necessary.    Filter  the 
solution  into   a  test- 
tube  containing  a 
clear  solution  of 
HgCl2. 

White  prec., 
(HgCl)2,  or  dark 
pree.,  Hg,  indicates 
Sn. 


FIG.  i. 


FIG.  2. 


FIG.  3. 


tested  in  a  platinum  crucible  cover  or  small  crucible  made  with  a  projecting  edge  for 
holding  with  pincers  (Fig.  2),  the  contact  may  be  easily  made  as  shown  in  Fig.  3.  The 
Sb  is  deposited  on  the  inside  of  the  Pt  crucible  which  thus  becomes  the  negative  elec- 
trode. 

J  Although  a  heavy  black  deposit  on  the  platinum  always  indicates  Sb,  a  light  stain 
may  be  due  to  Cu  if  this  element  is  present  in  the  original  solution.  CuS  is  somewhat 
soluble  in  large  amounts  of  yellow  ammonium  sulfid  and,  if  dissolved,  will  appear  on 
the  platinum  at  this  point. 

If  the  stain  is  light  and  Cu  is  known  to  be  present  in  the  original  solution,  dis- 
solve the  stain  by  warming  with  a  few  drops  of  HNO3  and  make  the  solution  alkaline 
with  NH4OH.  If  a  blue  color  does  not  develop,  the  stain  indicates  a  small  amount  of 
Sb  present.  If  a  blue  color  does  appear,  make  faintly  acid  with  HC1  and  pass  H2S. 
Treat  the  precipitated  sulfids  with  Na2S,  filter  and  acidify  with  HC1.  Orange  prec., 
Sb2SB,  indicates  Sb. 


THE    BASIC    ANALYSIS.  257 

DISCUSSION  OF  THE  ANALYSIS  OF  GROUP  I,  DIV.  C. 

The  separation  of  As  (and  of  Au  and  Pt,  if  present),  from  the  other  mem- 
bers of  this  Div.  depends  on  the  insolubility  of  the  sulfids  of  these  elements  in 
HC1  and  the  solubility  of  Sb2S6  and  SnS2.  The  acid  must  be  fairly  strong  in 
order  to  remove  completely  Sb  and  Sn,  yet,  since  boiling  is  necessary  to  remove 
H2S,  long  treatment  with  cone.  HC1  is  likely  to  dissolve  a  portion  of  the  As,  also. 
The  strength  prescribed,  about  8  N  HC1,  should  be  carefully  followed,  therefore, 
in  order  to  obtain  satisfactory  results. 

The  separation  of  Sb  from  Sn  depends  upon  the  different  action  of  HC1  upon 
the  two  metals.  By  bringing  into  contact  Pt  and  Zn  and  placing  them  in  a 
solution  of  a  salt,  a  difference  of  potential  results  and  an  electric  current  is  said 
to  flow.  Any  metallic  ions  in  solution  travel  in  the  direction  of  the  positive 
electricity  and  will  be  deposited  on  the  Pt  giving  up  their  charges  and  passing 
into  the  metallic  condition.  Antimony  is  only  slowly  attacked  by  hot  cone.  HC1 
while  tin  is  readily  acted  upon.  Hence  Sn  dissolves  as  rapidly  as  it  is  deposited 
and  Sb  alone  remains.  If,  however,  the  concentration  of  the  acid  falls,  Sn  may 
be  deposited  as  a  gray  crystalline  mass  either  on  the  Pt  because  of  the  electric 
current  or  on  the  Zn  because  of  the  greater  tendency  of  this  metal  to  pass  into 
the  ionic  condition.  The  smooth  black  deposit  of  Sb  is  not  likely  to  be  confused 
with  the  gray  crystals  of  Sn.  The  latter  will  dissolve  in  a  few  moments  on  adding 
cone.  HC1,  leaving  the  Sb  indication.  A  slight  stain  of  Sb,  if  moist  and  exposed 
to  the  air,  may  disappear  in  a  short  time  owing  to  oxidation. 

In  another  portion  of  Filt.  10,  Sn  may  be  detected  by  reducing  it  to  the 
stannous  condition.  The  color  of  the  solution  offers  some  indication  of  the  prog- 
ress of  the  reaction.  A  yellow  color  is  due  to  ferric  salts;  ferrous  salts  are 
nearly  colorless.  Since  stannic  salts  are  somewhat  more  difficult  to  reduce  than 
ferric  salts,  it  follows  that  the  reduction  is  not  complete  as  long  as  the  solution 
shows  any  yellow  color.  Aside  from  stannous  salts,  no  substances  able  to  reduce 
mercuric  chlorid  to  the  lower  chlorid  or  to  the  metal  will  be  present  in  this  solu- 
tion. Stannous  chlorid  has  already  been  used  as  a  reagent  for  detecting  the  pres- 
ence of  mercury.  Mercuric  chlorid  is  now  used  as  a  reagent  for  stannous  salts 
and,  if  a  reduction  takes  place  on  adding  HgCl2,  it  indicates  the  presence  of  Sn. 

Both  PtS2  and  Au2S3  are  almost  black  in  color,  hence  if  Res.  10  is  not  dark, 
neither  Pt  nor  Au  can  be  present.  It  is  only  necessary  in  this  case  to  oxidize  As2S5 
to  H3As04  (see  page  50).  Arsenic  compounds  in  general  are  easily  oxidized  but 
the  S  which  is  always  present  in  excess  from  the  decomposition  of  the  polysulfid 
renders  the  oxidation  in  this  case  somewhat  more  diffcult  by  mechanically  protect- 
ing the  As  from  the  action  of  the  acid.  Hence  it  is  necessary  to  evaporate  to  dry- 
ness  once  or  twice  to  insure  oxidation.  With  the  Ag  ion  the  As04  ion  produces  a 
chocolate-colored  precipitate  which  is  extremely  soluble  in  acid  and  in  NH4OH. 
Hence  it  can  be  precipitated  only  from  distinctly  neutral  solutions.  This  is 
readily  accomplished  by  floating  on  the  top  of  the  slightly  acid  solution  a  little 
dilute  NH4OH.  A  neutral  zone  will  form  between  the  two  layers  and  on  the  line 
of  contact  the  colored  precipitate  will  appear. 

Unless  Res.  10  is  washed  free  from  HC1,  the  test  for  As  may  be  obscured  by 
the  precipitation  of  AgCl.  If  such  is  the  case,  make  the  solution  slightly  acid  with 
HN03  and  filter  until  clear.  Upon  the  acid  nit.  float  NH4OH,  adding  more  AgNO, 
if  necessary,  and  look  for  the  indication  of  As  on  the  line  of  contact. 

18 


258 


QUALITATIVE    ANALYSIS. 


ANALYSIS  OF  RESIDUE  10  WHEN  GOLD  OR  PLATINUM  MAY 

BE  PRESENT. 

B.  //  Res.  10,  from  page  256,  is  dark,  i.  e.,  probably  contains  Au  or  Pt.  Punc- 
ture the  filter  and  wash  into  a  test-tube  with  10-15  ccm.  (NHJ2C03.  Shake  vigor- 
ously for  a  few  minutes,  filter  and  wash. 


Au2S3,  PtS2. 


RES.  11. 

Dissolve  in  as  little  aqua  regia  as  pos- 
sible and  evaporate  to  dryness  to  remove 
excess  of  acid.  Boil  out  with  15  ccm.  water 
and  divide  into  three  portions. 


A  (1)  Add  a  few 
drops  of  a  freshly  pre- 
pared FeS04  solution. 

Dark  prec.,  metallic 
An,  indicates  Au. 

A  (2)  Dilute  to  50 
ccm.    Add  a  single  drop 
of  SnCl4  and  a  drop  or 
two  of  SnCl2. 

Purple  or  brown  color, 
metallic  Au,  indicates 
Au. 


B.  Add 
5-10  drops  of 
KC1  and  a 
volume  of 
alcohol  twice 
as  great  as  the 
solution. 

Yellow 


prec., 
indicates  Pt. 


FILT.  11.     Make  faintly  acid 
with  HN03,  pass  H,S  if  necessary 
to  obtain  a  prec.,  filter  and  wash. 


PREC.  12. 

As,S, 

Yellow  prec., 
indicates  As. 
To  confirm, 
treat  as  Res. 
10  as  directed 
under  A, 
256. 


FILT.  12. 
Reject. 


THE    BASIC    ANALYSIS.  259 

DISCUSSION  OF  THE  ANALYSIS  OF  RESIDUE  10  WHEN  GOLD  OR 
PLATINUM  MAY  BE  PRESENT. 

If  either  Au  or  Pt  be  present  in  appreciable  amounts,  Res.  10  will  be  dis- 
tinctly dark  in  color,  for  the  sulfids  of  these  metals  are  black.  A  slight  darkening 
of  the  normal  yellow  color  of  Res.  10,  however,  may  be  due  to  sulfur  which  sepa- 
rates when  the  ammonium  polysulfid  extract  is  acidified  and  which,  under  these 
conditions,  is  often  somewhat  discolored. 

The  separation  of  As  from  Au  and  Pt  depends  upon  the  solubility  of  the. 
sulfids  of  arsenic  in  (NH4)2C03,  according  to  the  equation: 

AsA  +  3(NH4),CO,-»  (NH4)3AsS4+  (NH4)3AsS03  +  3C02 

Ammonium  thioarsenate  and  the  arsenate  in  which  only  a  quarter  of  the 
oxygen  has  been  replaced  by  sulfur  both  decompose,  when  acidified,  with  the 
formation  of  As2S5  which  may  be  confirmed  by  the  previous  tests.  See  page  256. 

The  action  of  aqua  regia  on  the  sulfids  of  gold  and  platinum  converts  them 
into  soluble  chlorids.  The  tests  for  gold  depend  upon  the  ready  reduction  of  its 
salts  to  the  metallic  condition.  Ferrous  salts  precipitate  gold  according  to  the 
equation : 

2AuCl3  +  6FeS04  -»  2Au  +  2FeCl3  +  2Fe2(S04)3 

Similarly,  stannous  salts  precipitate  gold  in  a  very  finely-divided  condition, 
according  to  the  equation : 

2AuCl3  +  3SnCl2  -»  2Au  +  3SnCl4 

This  precipitate,  known  as  the  "  Purple  of  Cassius,"  always  contains  some 
stannous  oxid,  SnO. 

The  platinic  chlorid,  PtCl4,  formed  by  the  action  of  aqua  regia  upon  PtS2, 
unites  immediately  with  the  HC1  present  in  the  aqua  regia  mixture  to  form  hydro- 
chlorplatinic  acid,  H2PtCl6.  The  potassium  salt  of  this  acid  is  somewhat  soluble 
in  water  but  much  less  so  in  alcohol  and,  on  the  addition  of  the  latter  reagent,  the 
yellow  precipitate,  K2PtCl6,  falls,  indicating  Pt.  (See  pages  74  and  235.) 


260 


QUALITATIVE    ANALYSIS. 


DETECTION  OF  INTERFERING  SUBSTANCES. 

Evaporate  a  small  portion  of  Filt.  4  to  dryness  in  a  small  evaporating  dish 
and  desiccate  at  110°-120°.    Designate  this  dried  res.  as  Prec.  13.     Divide. 

remainder  of  Prec.  13  with  HC1  and  filter. 
FILT.  14.     Divide  into  two  portions. 


A.  If  an  in- 

Boil out  the 

dication  has  not 

B.  PREC.  14. 

been  given  dur- 

Heat before 

ing  the  desicca- 

the blow- 

tion, heat  in 

pipe  in  a 

porcelain  a 

bead  of 

portion  of 

NaP03.* 

Prec.  13  with 

An  insolu- 

cone. H2S04 

ble  skeleton, 

until  white 

Si02, 

fumes  of  S08 

indicates 

appear. 

H2SiOs. 

Carboniza- 

tion, evidenced 

by  darkening 

and  "  odor  of 

burned  sugar," 

indicates  car- 

bonaceous 

matter. 

C.  To  the  smaller  portion 
of  Filt.  14  add  a  few  drops 
of  cone.  HN03  and  evaporate 
to  dryness.  Boil  out  the  res. 
with  5-10  ccm.  of  water  to 
which  a  drop  or  two  of  HN03 
has  been  added. 

To  5  ccm.  of  (NHJ2Mo04 
solution,  add  a  few  drops  of 
the  extracted  solution  and 
warm  slightly. 

Yellow  prec.,  (NH4)3P04. 
12MoO3,  indicates  H3P04. 

If  a  yellow  prec.  does  not 
form,  add  1-2  ccm.  of  the  ex- 
tracted solution  to  the  molyb- 
date  and  warm. 

If  no  prec.  appears,  test  all 
of  the  extract  in  the  same 
way. 


D.  To  the 
larger  portion 
of  Filt.  14  add 
solid  Na2C03  to 
alkaline  reac- 
tion and  some- 
what in  excess, 
then  boil  thor- 
oughly.    Filter, 
acidify  the  filt. 
slightly  with 
H(C2H302) 
and  add  an 
equal  volume  of 
saturated 
CaS04  solution. 

White    prec., 
CaC204,  indi- 
cates H2C204. 


DISCUSSION  OF  THE  DETECTION  OF  INTERFERING  SUBSTANCES. 
Certain  substances  interfere  with  the  separation  and  detection  of  some  of  the 
metals.  So  long  as  the  solution  is  acid  with  HC1  and  a  considerable  concentration 
of  H  ions  is  present,  the  reactions  noted  below  do  not  take  place.  On  making  the 
solution  ammoniacal  or  faintly  acid  with  acetic  acid,  these  substances  begin  to 
interfere  and,  therefore,  must  be  removed  before  proceeding  further  with  the 
analysis. 

Organic  bodies  containing  the  hydroxyl  radicle,  such  as  sugar,  starch,  tartaric, 
citric  and  other  acids,  prevent  the  precipitation  of  ferric  iron,  aluminum  and,  to 
a  smaller  extent,  of  base-forming  chromium  as  hydroxids  or  basic  acetates.  Com- 
plex bodies  are  formed  which  do  not  separate  the  simple  metallic  ions  to  an  extent 
sufficient  to  cause  precipitation. 
!  If  silicic  acid,  H2Si08,  is  present,  ammonium  silicate  will  begin  to  form  as 

*  Sodium  meta-phosphate,  NaPO3,  is  usually  prepared  in  small  amounts  as  needed 
by  the  ignition  of  microcosmic  salt,  H(NH4)NaPO4.  A  small  loop  of  platinum  wire 
is  heated  and  brought  in  contact  with  the  microscosmic  salt  crystals  which  adhere  to 
it.  On  heating,  NH3  and  H2O  are  driven  off,  leaving  a  clear,  colorless  bead  of  NaPO3. 
By  pressing  the  hot  bead  upon  the  substance  to  be  tested,  a  sufficient  quantity  will 
adhere  to  give  the  reaction  on  heating.  The  skeleton  of  silica  is  more  noticeable  during 
the  heating  than  subsequently,  since  it  spins  about  in  the  bead  under  the  influence  of 
the  blowpipe  flame.  (See  page  78.) 


THE    BASIC    ANALYSIS.  261 

soon  as  the  solution  is  made  nearly  or  quite  alkaline  with  (NH4)2C03  or  NH4OH 
as  in  the  analysis  of  Groups  II-IV.  Ammonium  hydroxid  is  a  weak  base  and 
silicic  acid  one  of  the  weakest  of  acids,  hence  the  salt  is  completely  hydrolyzed  and 
separates  silicic  acid  as  a  white  gelatinous  precipitate  which  is  likely  to  be  mistaken 
for  aluminum  hydroxid  or  other  white  precipitate. 

Phosphoric  and  oxalic  acids  form  insoluble  salts  with  every  metal  still  present 
in  solution  with  the  exception  of  the  alkalies.  Some  of  these  compounds  will  pre- 
cipitate in  the  presence  of  acetic  acid  but  all  will  fall  as  soon  as  the  solution  is 
made  ammoniacal.  Unless  these  acids  are  removed,  therefore,  members  of  Groups 
II-IV  may  be  precipitated  in  the  wrong  place  in  the  analysis  and  thus  introduce 
confusion. 

A.  Most  non-volatile  organic  substances  decompose  when  heated  with  the  loss 
of  water  and  the  separation  of  free  carbon.    This  "  carbonization  "  or  "  charring  " 
may  be  noted  during  the  preliminary  evaporation  to  dryness  inasmuch  as  the 
sides  of  the  dish  above  the  liquid  usually  become  heated  sufficiently  high  to  cause 
decomposition.    Hot  concentrated  H2S04  tends  to  extract  water  and  thus  aids  the 
decomposition. 

B.  Silicic  acid,  H2Si03,  if  present  in  solution,  becomes  dehydrated  during  the 
desiccation,  forming  the  insoluble  anhydrid,  Si02,  which  constitutes  the  residue 
left  after  extracting  with  HC1.    Heated  in  a  bead  of  NaP03,  no  reaction  takes 
place.     Sodium  meta-phosphate  dissolves  base-forming  oxids  with  the  formation 
of  ortho-phosphates  which  fuse  to  a  clear  glass  but  non-volatile,  infusible,  acid- 
forming  oxids,  like  Si02,  remain  undissolved  as  an  opaque  residue.     (See  page 
136.) 

C.  In  the  present  scheme  of  analysis,  compounds  containing  phosphorus  are 
separated   and  detected  only  after  the   phosphorus   has  been   oxidized   to  the 
pentavalent  condition  and  exists  as  ortho-phosphoric  acid,  H3P04,  for  other  forms 
do  not  respond  to  the  reactions  given.     Evaporation  with  cone.  HN03  accom- 
plishes the  oxidation  and  hydration  is  brought  about  by  boiling  with  the  dilute  acid. 

The  test  for  H3P04  depends  upon  the  fact  that  in  the  presence  of  a  large 
excess  of  molybdic  acid,  H2Mo04,  phosphoric  acid  unites  with  the  anhydrid,  Mo03 

to  form  a  complex  acid.  In  the  presence  of  NH4  ions,  ammonium  phospho- 
molybdate,  (NH4)3P04.12Mo03,  is  formed,  a  compound  insoluble  in  HN03.  This 
yellow  precipitate  will  come  down  slowly  in  the  cold  but  slight  warming  will 
hasten  its  formation.  Even  moderate  heating  is  to  be  avoided,  however,  as  it  tends 
to  dehydrate  molybdic  acid  and  separate  the  straw-colored  anhydrid.  It  is  neces- 
sary that  the  proportions  of  Mo  to  P  be  at  least  12 :1  for  otherwise  compounds 
containing  less  Mo  will  form  which  are  soluble.  Hence  the  necessity  of  adding 
only  a  drop  of  the  solution  to  be  tested  to  5  ccm.  of  (NH4)2Mo04  in  order  to 
prevent  the  possibility  of  an  excess  of  H3P04,  even  should  it  be  present  in  con- 
siderable amounts  in  the  solution.  The  test  is  extremely  delicate  and  serves  to 
detect  traces  of  phosphoric  acid. 

D.  The  addition  of  Na2C03  to  the  filtrate  precipitates  all  the  metals  present  in 
solution  either  as  carbonates  or  as  insoluble  salts  of  other  acids  present.     On 
boiling  the  mixture,  however,  the  latter  are  all  transposed  into  carbonates  of  the 
metals  and  soluble  sodium  salts  of  the  acids  present.     By  this  treatment,  there- 
fore, all  bases  except  the  alkalies  are  removed  while  the  acids  remain  in  solution. 
The  addition  of  a  solution  of  CaS04  (soluble  in  500  parts  of  H20)  introduces 
Ca  ions  in  such  concentration  that  only  highly  insoluble  calcium  salts  can  be  pre- 
cipitated.    Acetic  acid  prevents  the  formation  of  the  phosphate,  carbonate  and 
similar  salts  and  only  the  oxalate  (soluble  in  140,000  parts  of  H20  and  insoluble 
in  acetic  acid)  can  precipitate. 


262  QUALITATIVE    ANALYSIS. 

REMOVAL    OF    INTERFERING    SUBBSTANCES. 

A.  If  no  interfering  substances  are  detected,  proceed  with  the  analysis  of 
Filt.  4  as  directed  on  page  264. 

B.  If  carbonaceous  matter  or  oxalic  acid,  H2C,04,  or  both  are  detected,  evapor- 
ate to  dryness.     Grind  the  dried  residue  in  a  mortar  with  4-5  times  its  bulk  of 
solid  NH4N03.    In  a  porcelain  crucible  fuse  enough  NH4N03  to  cover  the  bottom 
and  add  the  mixture  in  small  portions,  so  regulating  the  temperature  as  to  keep 
the  contents  of  the  crucible  in  a  state  of  gentle  fusion.    When  quiet  fusion  of  the 
melt  indicates  that  the  oxidation  is  complete,  raise  the  temperature  somewhat  to 
drive  off  most  of  the  NH4N03,  cool  and  extract  with  hot  water,  adding  a  little  HC1 
if  necessary.    Designate  this  solution  as  Filt.  15.    If  it  is  colorless  (i.  e.,  does  not 
contain  H2Cr2Or)  and  is  free  from  HJSi03,  proceed  according  to  the  directions 
for  the  analysis  of  Filt.  4  on  page  264.    If  Filt.  15  is  somewhat  yellow  in  color  and 
a  small  portion  when  tested  according  to  A  (2),  page  266,  indicates  H2Cr207,  con- 
centrate to  5-10  ccm.    Add  an  equal  volume  of  cone.  HC1  and  boil  to  reduce  the 
chromium  from  the  acid-forming  to  the  base-forming  condition.    The  reduction  is 
indicated  by  a  change  in  color  from  yellow  to  deep  green.     The  completion  of 
the  reaction  may  be  hastened  by  the  addition  of  a  few  drops  of  alcohol.    Remove 
most  of  the  HC1  by  evaporating  nearly  to  dryness  and  dissolve  in  water.    Desig- 
nate this  solution  as  Filt.  16  and  (if  free  from  H2Si03)  proceed  according  to  the 
directions  for  the  analysis  of  Filt.  4  on  page  264. 

C.  If  silicic  acid,  H2SiO3,  is  indicated,  evaporate  to  dryness  in  the  presence 
of  HC1  and  desiccate  at  110-120°.    Boil  out  with  water  and  a  little  HC1  and  care- 
fully decant  the  liquid  from  the  insoluble  residue,  pouring  it  through  a  filter. 
Wash  the  residue  by  decantation  with  hot  water  and  pour  the  washings  through 
the  filter. 


RES.  17.  Dry  and  transfer  to  a  platinum  crucible. 
Add  a  few  drops  of  HF  and  about  1  ccm.  of  cone.  H2S04 
and  heat  under  a  hood  until  fumes  of  S03  are  no  longer 
given  off.  If  an  appreciable  residue  is  left,  boil  out  with 
HC1  and  pour  off  the  solution  through  a  filter. 

RES.  18.  If  appreciable,  fuse  in  platinum  with  3-4 
times  its  bulk  of  a  mixture  of  Na2C03  and  K2C03.  After 
the  melt  becomes  clear  and  fuses  quietly,  cool  and  boil 
out  with  water.  Filter. 


FILT.  17.    After  the 
addition  of  Filt  18, 
proceed  as  directed 
under  Filt.  4,  page 
264. 

FILT.  18.    Add  to 
Filt.  17. 


RES.  19.     Wash  with  hot  water,  extract  with 
HC1  and  filter. 


FILT.  19.     Reject. 


RES.  20.     If  any,  reject.  FILT.  20.     Test  for  members  of 

Group  IV  as  directed  under  Filt. 
23,  page  270. 

D.  Phosphoric  acid,  H3P04,  if  present,  is  removed  in  the  separation  of  Group 
II  as  directed  on  page  264. 


THE    BASIC   ANALYSIS.  263 

DISCUSSION  OF  THE  REMOVAL  OF  INTERFERING  SUBSTANCES. 

B.  The  removal  of  oxalic  acid  and  of  carbonaceous  matter  in  general  depends 
upon  the  oxidation  of  these  substances  to  C02  and  H20.    Nitrates  are  very  satis- 
factory oxidizing  agents,  under  these  conditions  the  NH4  rather  than  the  Na  or  K 
salt  being  preferable,  since  it  introduces  no  metallic  element  into  the  solution. 
(Tests  are  made  for  the  presence  of  NH4  salts  before  beginning  the  systematic 
analysis.    See  page  249.)    During  this  fusion,  Cr  (if  present)  will  be  oxidized  into 
the  higher  or   acid-forming  condition.      Since  the   separation   of  this   element 
depends  on  its  being  in  the  base-forming  condition  it  must  be  reduced.     This 
is  accomplished  by  boiling  with  cone.  HC1,  chlorin  being  liberated  according  to  the 
reaction : 

H2Cr207  +  12HC1  ->  2CrCl3  +  7H20  +  3C12 

In  the  presence  of  the  strong  oxidizing  agent,  dichromic  acid,  HC1  plays  the 
part  of  a  reducing  agent  (see  page  49)  but  it  is  not  readily  oxidized  and  boiling  for 
some  minutes  is  necessary  to  complete  the  reaction.  Alcohol,  a  strong  reducing 
agent,  is  usually  added  to  hasten  the  action  (see  pages  47  and  192)  and  is 
oxidized  according  to  the  equation : 

H2Cr207  +  6HC1  +  3H6C20  -»  2CrCl3  +  3H4<720  +  7H20 

The  product  of  the  reaction,  aldehyde,  H4C20,  being  volatile,  may  be  removed 
by  boiling  from  the  solution,  together  with  any  unchanged  alcohol. 

If  silicic  acid  be  present  during  this  fusion,  insoluble  silicates  of  some  of  the 
metals  present  may  be  formed  which  will  not  be  decomposed  by  HC1  and  there- 
fore require  the  treatment  indicated  under  Res.  17. 

C.  Silicic  acid  is  removed  from  solution  by  drying  or  desiccating  at  110-120°, 
at  which  temperature  it  dehydrates  forming  the  insoluble  anhydride,  Si02.     If 
the  desiccation  is  properly  performed,  silicic  acid  may  be  removed  so  completely 
that  it  will  not  interfere  with  the  further  analysis  of  the  solution.    If,  however, 
the  desiccation  is  imperfect,  silicic  acid  may  pass  into  Filt.  17,  in  which  case 
this  must  be  evaporated  and  desiccated  a  second  time.    Heating  higher  than  120° 
must  be  avoided,  for,  above  this  temperature,  the  SiO,  first  formed  may  react  with 
metallic  chlorids  present  to  form  soluble  silicates  and  thus  introduce  silicic  acid 
into  the  filtrate. 

2NaCl  -f  Si02  +  H20  -»  Na,Si03  +  2HCI 

A  very  satisfactory  way  to  maintain  the  temperature  necessary  for  desiccation 
and  yet  prevent  overheating  is  to  fan  the  evaporating  dish  with  a  Bunsen  flame, 
stirring  the  drying  residue  constantly  with  a  glass  rod  and  breaking  up  any  lumps 
that  tend  to  form,  until  a  fine  dry  powder  is  obtained. 

Silicic  anhydrid  and  all  silicates  are  decomposed  by  HF  with  the  formation  of 
volatile  SiF4  (see  page  135).  On  heating  Res.  17  with  HF  and  H2S04,  all  Si  will 
be  driven  off  and,  if  the  analysis  has  been  properly  performed  up  to  this  point, 
there  will  usually  be  only  an  inappreciable  amount  of  material  left  as  Res.  18. 
The  sulfates  of  Ba,  Sr  and  Ca  may  sometimes  be  present  in  considerable  quantities, 
however,  having  been  formed,  either  by  incomplete  removal  of  H2S  from  Filt. 
4  (see  page  253)  and  its  subsequent  oxidation  into  H2S04,  or  by  the  decomposition 
by  means  of  HF  and  H2S04  of  insoluble  silicates  of  these  same  metals,  the  silicates 
having  been  formed,  possibly,  during  the  fusion  as  indicated  in  B.  If  this  be  the 
case,  fusion  with  an  alkaline  carbonate  transposes  them  into  carbonates  of  Ba, 
Sr  and  Ca  and  sulfates  of  the  alkali  metal  (see  page  69).  Extraction  with  water 
removes  the  latter  after  which  the  carbonates  may  be  dissolved  by  treating 
with  HC1. 


264  QUALITATIVE    ANALYSIS. 

SEPARATION  OF  GROUP  II. 
To  filt.  4,  freed  from  H2S,  add  a  little  HN03,  boil  and  filter,  if  necessary. 


RES.  21.*  BaS04,  SrS04,  CaSO4. 

Treat  as  directed  for  Res.  18  on  page  262. 


FILT.  21.     Divide  into  two 
unequal  portions. 


A.  To  a  small  portion  add  2-3  ccm.  NH4C1,  if  it  is  not  already  present  in 
sufficient  amount,f  make  faintly  alkaline  with  NH4OH  and  boil.    A  prec.  indicates 
the  presence  of  Group  II  (or  Groups  III  or  IV  as  phosphates).     //  no  prec. 
forms,  proceed  with  the  separation  and  analysis  of  Group  III  as  directed  on 
page  268. 

B.  //  a  prec.  forms  in  A,  add  5-10  ccm.  NH4C1,  if  it  is  not  already  present  in 
sufficient  amount.f    Add  (NH4)2C03  until  the  first  formation  of  a  prec.  indicates 
the  neutralization  of  the  solution.    If  the  prec.  which  falls  is  not  colored  reddish- 
brown,  add  FeCl3  until  such  a  prec.  does  form  on  the  further  addition  of  a  few 
drops  of  (NH4)2C03.    Prevent  the  formation  of  any  large  amount  of  prec.  by  dis- 
solving in  HC1  after  each  addition  of  (NH4)2C03  if  the  prec.  is  not  of  the  right 
color.    When  sufficient  FeCl3  has  been  added,  dissolve  the  prec.  in  the  smallest 
possible  amount  of  HC1  and  add  10  ccm.  of  NH4(C2H302).     Pour  the  resulting 
deep  red  solution  into  400-500  ccm.  of  boiling  water  and  maintain  the  boiling 
temperature  until  a  reddish-brown  prec.  forms.    Filter  hot. 


PREC.  22.  Group  II.  Al,  Or  and  Fe  as  basic  acetates. 
(FeP04,  if  H3P04  is  present.)  Proceed  as  directed  on 
page  266. 


FILT.  224  Groups 
III-V.    See  page 
268. 


*  This  res.,  if  any,  is  due  to  the  incomplete  removal  of  H2S  and  its  subsequent 
oxidation  to  H2SO4.  If  the  analysis  has  been  performed  carefully  as  directed,  there 
should  ordinarily  be  no  Ees.  21. 

t  Sufficient  NH4C1  will  ordinarily  be  formed  during  the  course  of  an  analysis  from 
the  neutralization  of  HC1  with  NH4OH  or  (NH4)2CO3. 

|  Filt.  22  should  be  free  from  Fe  (and  Al  and  Cr  also)  and  when  acidified  with 
HC1,  should  give  no  test  with  K4(Fe  =  CGN6)  (see  page  227).  If  Fe  should  be  indi- 
cated, the  separation  has  not  been  satisfactorily  performed,  owing,  probably,  to  the 
presence  of  too  much  acetic  acid,  H(C2H3O2).  Under  these  conditions  the  filtrate  should 
be  heated  to  boiling  and  a  few  drops  of  (NH4)2CO3  added,  not  sufficient,  however,  to 
make  the  solution  neutral  or  alkaline.  Filter  off  the  prec.  and  add  to  Prec.  22. 


THE    BASIC    ANALYSIS.  265 

DISCUSSION  OF  THE  SEPARATION  OF  GROUP  II. 

The  separation  of  Group  II  depends  on  the  fact  that  aluminum,  chromium, 
and  iron  form  basic  acetates  which  are  insoluble  in  the  presence  of  a  small  con- 
centration of  acetic  acid,  whereas  the  metals  of  Groups  III-V  do  not  possess  this 
property.  For  the  formation  of  these  compounds  it  is  necessary  that  chromium 
be  present  in  the  trivalent  base-forming  condition  and  not  in  the  higher  oxidized 
acid-forming  condition.  Since  ferrous  iron  does  not  form  a  basic  acetate,  it  is 
necessary  that  iron  be  in  the  trivalent  condition.  If  chromium  be  present  in  the 
original  solution  in  the  acid-forming  condition  and  iron  as  a  ferric  salt,  both  will 
be  reduced  by  H2S  in  the  presence  of  HC1,  chromium  to  the  trivalent  and  iron  to 
the  bivalent  condition.  In  Filt.  4  it  is  necessary,  therefore,  to  oxidize  the  iron 
without  oxidizing  the  chromium.  Nitric  acid  in  dilute  solution  accomplishes  this 
satisfactorily.  If  H2S  has  not  been  previously  boiled  out  completely  (see  page 
252),  it  will  be  oxidized  to  H2S04  by  the  HN03  and  cause  the  precipitation  of  the 
metals  of  Group  IV  (Res.  21). 

Inasmuch  as  a  preliminary  test  may  indicate  the  complete  absence  of  Group 
II,  it  is  advantageous  to  test  Filt.  21  for  the  presence  of  this  group  before  pro- 
ceeding to  the  regular  separation.  The  test  as  directed  in  A  depends  upon  the 
fact  that  the  hydroxids  of  aluminum,  chromium  and  ferric  iron  are  so  little  solu- 
ble in  water  that  they  may  be  precipitated  by  NH4OH  even  in  the  presence  of 
salts  of  ammonium  which  repress  the  ionization  of  the  base  (see  page  31).  The 
hydroxids  of  the  metals  of  Groups  III-V,  being  more  soluble  in  water,  cannot  be 
precipitated  by  NH4OH  when  the  concentration  of  hydroxyl  ions  is  reduced  by 
the  presence  of  ammonium  salts.  If  H3P04  is  present,  however,  Groups  III  and  IV 
may  be  precipitated  as  phosphates,  since  the  phosphates  of  all  of  the  metals, 
though  soluble  in  acids,  are  insoluble  in  water  and  are  precipitated  as  soon  as  the 
solution  is  made  alkaline.  Inasmuch  as  the  basic  acetate  separation  removes  H3P04 
as  well  as  the  metals  of  Group  II,  if  a  precipitate  forms  in  A,  the  regular  pro- 
cedure for  the  separation  and  analysis  of  Group  II  must  be  followed. 

In  order  to  make  this  separation  satisfactorily,  the  solution  should  contain 
as  little  acid  as  possible  in  excess  of  the  amount  necessary  to  give  an  acid  re- 
action. If  the  precipitate  formed  by  (NHJ2C03  in  B  is  dissolved  in  the  smallest 
possible  amount  of  HC1,  on  addition  of  NH4(C2H302),  acetic  acid  (and  NH4C1) 
will  be  formed  in  place  of  the  stronger  HC1  and  on  dilution  to  500  ccm.  the  con- 
centration of  H  ions  will  be  sufficiently  low  to  allow  the  acetates  to  hydrolyze  and 
precipitate  the  insoluble  basic  salts.  To  what  extent  hydrolysis  of  the  normal 
acetates  of  the  type  Fe(C2H302)3  will  take  place  depends  somewhat  upon  con- 
ditions. The  composition  of  the  precipitated  basic  acetates  varies  from 
Fe(OH)(C2H302)2  to  Fe(OH)2(C2H3OJ.  The  filtration  must  be  made  while 
the  solution  is  hot  to  prevent  a  reversal  of  the  reaction  and  the  solution  of  the 
basic  acetates  in  the  acetic  acid  formed  by  hydrolysis. 

If  H3P04  be  present,  FeP04  will  form  in  preference  to  any  other  salt,  since 
it  is  least  soluble  in  water  and  acetic  acid.  This  is  indicated  by  the  fact  that  on 
the  addition  of  FeCl3  to  a  solution  containing  phosphoric  acid,  straw-colored 
ferric  phosphate  is  precipitated  on  neutralizing  the  acid  rather  than  the  red 
hydroxid.  Only  after  all  the  H3P04  present  has  united  with  the  iron,  does  the 
hydroxid  fall.  That  the  excess  of  iron  necessary  for  complete  precipitation  of 
chromium  (see  page  194)  is  present  is  indicated  by  the  appearance  of  the  red 
precipitate  in  B. 


266 


QUALITATIVE    ANALYSIS. 


ANALYSIS  OF  GROUP  II. 

PREC.  22,  from  page  264.     Group  II.    Al,  Cr,  and  Fe*  as  basic  acetates. 
(FeP04,  if  H3P04  is  present.)     Divide  into  two  portions. 


A.  Mix  one  portion  of  the  pree.  with 
5-10  ccm.  water  and  add  a  little  Na,02, 
stirring  or  shaking  the  mixture  to  in- 
sure thorough  treatment.  Filter. 

A  yellow  nit.  indicates  Cr. 

Confirm  by — 

(1)  Acidifying  carefully  a  portion 
of  the  filt.  with  H2S04.     A  transient 
blue   color,   HCr04,   appearing  as  the 
solution  becomes  acid  or  on  adding  a 
few  grains  of  Na,0,,  confirms  Cr. 

(2)  Evaporating  a  second  portion  of 
the  filt.  to  dryness,  cool,  boil  out  with 
water,  faintly  acidify  with  H(C2H302) 
and  addPb  (C2H302)2. 

Yellow  prec.,  PbCr04,  confirms  Cr. 


B.  Mix  the  second  portion  of  the 
prec.  with  5-10  ccm.  water,  add  about 
half  a  gram  of  solid  Na,C03  and  nearly 
as  much  solid  Ba(OH)2  and  boil  thor- 
oughly. Filter  into  a  test-tube,  faintly 
acidify  with  HC1,  make  just  alkaline 
with  NH4OH  and  boil  until  a  prec. 
forms  or  the  odor  of  NH3  disappears. 

Light  white  floccy  prec.,  A1(OH)3, 
more  readily  seen  against  a  dark  back- 
ground, indicates  Al. 

To  confirm,  filter  and  wash  the  prec. 
into  the  tip  of  the  paper.  Tear  off  the 
tip  of  the  filter  containing  the  prec. 
and  ignite  strongly  in  platinum.  Mois- 
ten with  drop  of  dilute  Co(N03)2  and 
ignite  again. 

Blue  color,  Co(A102)2,  confirms  Al. 


*  Special  tests  are  made  for  the  presence  and  condition  of  iron  before  beginning 
the  Basic  Analysis  (see  page  249).  Since  FeCl3  is  added  to  the  solution  to  insure  the 
complete  precipitation  of  Cr  (and  to  remove  H3PO4,  if  present),  it  would  obviously  be 
absurd  to  test  for  the  presence  of  Fe  in  this  group. 


UNIVERSITY 

/       THE    BASIC    ANALYSIS.  267 


DISCUSSION  OF  THE  ANALYSIS  OF  GROUP  II. 

A.  The  test  for  Cr  depends  on  the  oxidation  of  the  element  from  the  trivalent, 
base-forming  condition  in  which  it  is  precipitated  as  the  basic  acetate,  to  the 
hexavalent  acid-forming  condition.     Sodium  peroxid,  Na202,  reacts  with  water 
to  form  H,02  and  NaOH,  hence  it  is  equivalent  to  H202  in  alkaline  solution. 
Under  these  conditions  the  oxidation  of  Cr  stops  with  the  formation  of  an  alkali 
chromate,  e.  g.,  Na,Cr04,  which  gives  a  yellow  filtrate. 

Cr(OH)2.(C2H30,)  -f  NaOH  -»  Cr(OH)3  +  Na(C,H,<X) 
2Cr(OH)3  +  3H8Oa  +  4NaOH  ->  2Na2Cr04  +  8H20 

The  confirmatory  tests  depend  upon  two  reactions  of  chromium  in  the  acid- 
forming  condition.  In  acid  solution,  H202  oxidizes  hexavalent  Cr  one  step  higher 
than  in  alkaline  solution  and  forms  a  blue  compound,  probably  HCr04.  If  Na202 
was  not  added  in  sufficient  quantity  to  form  enough  H202  for  this  test,  it  may  be 
necessary  to  add  a  few  grains  after  acidifying.  A  large  excess  should  be  avoided, 
however,  because  HCr04  and  H202  interact  as  shown  by  the  equation: 

2HCr04  +  4H202  ->  Cr203  -f-  5H20  -f  402 

In  the  presence  of  acid,  Cr203  will  be  changed  into  the  corresponding  salt  of 
base-forming  Cr  which,  of  course,  will  not  give  a  blue  color. 

In  order  to  prevent  the  oxidation  of  Na2Cr04  by  H202  on  acidifying  with 
H(C,H302)  in  (2),  the  solution  is  evaporated  to  dryness  to  remove  all  H202,  after 
which  the  regular  test  with  the  Pb  ion  may  be  obtained. 

B.  The  separation  of  Al  from  Cr  depends  on  the  solubility  of  the  basic  acetate 
of  Al  in  boiling  NaOH  solution,  whereas  the  corresponding  salt  of  Cr,  though 
soluble  in  the  cold,  is  precipitated  on  boiling.    For  this  reaction,  the  NaOH  must 
be  prepared  by  the  action  of  Na2C03  on  Ba(OH)2;  for  the  ordinary  reagent,  espe- 
cially after  standing  in  glass  bottles,  always  contains  more  or  less  of  Al  or  com- 
pounds of  silicic  acid  which  may  be  confused  with  it.    When  made  by  the  action  of 
Na202  on  water,  the  same  impurities  are  present,  also,  and  are  likely  therefore, 
to  vitiate  the  test. 

In  the  filtrate  thus  obtained,  Al  is  present  as  sodium  aluminate,  Na3A103. 
Acidifying  with  HC1  and  making  alkaline  with  NH4OH  effects  the  change  from 
sodium  aluminate,  which  is  not  decomposed  by  boiling,  to  ammonium  aluminate, 
from  which  A1(OH)3  may  be  precipitated  on  heating.  (See  pages  116  and  117.) 


268 


QUALITATIVE    ANALYSIS. 


SEPARATION  AND  ANALYSIS  OF  GROUP  III. 

Heat  Filt.  22  nearly  to  boiling,  make  slightly  alkaline  with  NH4OH,*  add  2-3 
ccm.  colorless  (NH4)2S  and  filter.  Test  the  filt.  with  a  drop  of  (NH4)2S  to  ascer- 
tain that  precipitation  is  complete.  Wash  the  prec.  with  warm  water  to  which  a 
little  NH4C1  and  (NH4)2S  have  been  added. 


FILT.   23.f     Groups 
IV-V.     See  page  270. 


PREC.  23.     Group  III.    CoS,  NiS,  MnS,  ZnS.    Punc- 
ture the  filter  and  wash  the  prec.  into  a  test-tube  with 

10-20  ccm.  of  N  HC14  Shake  vigorously  for  5-10  minutes  to  insure  thorough 
treatment.  If  the  prec.  is  light  colored,  not  otherwise,  heat  may  be  applied  to 
hasten  solution.  Filter  and  wash. 

FILT.  24.  Boil  to  remove  H2S, 
make  alkaline  with  NaOH  and 
filter. 


RES.  24.  CoS,  NiS.  Heat  a  small  portion 
in  a  borax-bead.  Blue  color,  Co(B02)2,  indi- 
cates Co. 

A.  //  Co  is  not  indicated,  dissolve  the  pree. 
in  the  least  possible  amount  of  aqua  regia  and 
make  alkaline  with  Na2C03. 

Green  prec.,  NiC03.Ni(OH)2,  indicates  Ni. 

B.  //  Co  is  indicated,  dissolve  the  prec.  in 
the  least  possible  amount  of  aqua  regia,  evap- 
orate nearly  to  dryness  and  dilute  to  5  ccm. 
with  water.    Add  Na2C03  until  a  prec.  begins 
to  form  and  dissolve  it  again  in  the  least  pos- 
sible amount  of  H(C2Ha02).     Add  1-2  ccm. 
of  A1C13  or  alum  solution  and  a  small  crystal 
of   K3(Fe=C6N6)    and   shake   to   hasten   its 
solution.    Add  cone.  NH4OH  equal  to  half  the 
volume  of  the  solution,  shake  thoroughly  and 
filter. 


PREC.  25. 
Maroon  prec., 
Co3(Fe=C6N6)2, 

confirms  Co. 


PREC.  26.  Light- 
brown  prec., 
Mn(OH)2,      indi- 
cates Mn. 

To  confirm — 

A.  Dissolve      a 
small    portion    of 
the  prec.  in  10-20 
ccm.     HN03    and 
boil,  add  about  1 
gram     Pb02     and 
boil  again.    Stand 
to  allow  suspended 
matter    to    settle. 
Pink  color, 
HMn04,  indicates 
Mn. 

B.  Fuse  the 
rest   of  the   prec. 
with  Na2C03   and 
KN03  in  a  loop 
of     platinum 
wire. 

Green  color, 
Na2Mn04,     con- 
firms Mn. 


FILT.  26. 
Acidify 
faintly  with 
HC1,  add 
Na(C2H302) 
and  pass 
H2S 

White 
prec.,  ZnS, 
indicates 
Zn. 

Confirm 
by   igniting 
the  prec.  on 
platinum, 
moisten 
with  a  drop 
of  dil 
Co(N03)2 
and  ignite 
again. 

Green 
color, 
CoZnCX 


confirms 
Zn. 


FILT.  25.  To  the  filt.  add 
a  piece  of  solid  NaOH  the 
size  of  a  bean  and  stand 
aside. 

Black  prec.,  Ni(OH)3, 
forming  about  the  NaOH, 
indicates  Ni. 

If  no  prec.  appears  on 
standing,  boil  until  it  does 
form  or  the  odor  of  NH3  is 
gone. 

*  Should  a  prec.  form  on  the  addition  of  NH4OH,  sufficient  NH4C1  is  lacking. 
Dissolve  the  prec.  in  HC1  and  add  NH4C1  until  no  prec.  appears  when  the  solution  is 
again  made  alkaline.  If  an  analysis  of  Group  III  only  be  undertaken  without  the 
previous  treatment  to  separate  Groups  I  and  II,  15-20  ccm.  NH4C1  must  be  added  to 
the  sol.  before  the  addition  of  NH4OH.  In  the  course  of  a  systematic  analysis  this 
is  not  usually  necessary,  as  NH4C1  will  be  formed  in  sufficient  quantity  during  the 
previous  reactions. 

t  If  Filt.  23  is  brown,  it  is  due  to  NiS  which  is  somewhat  soluble  in  an  excess  of 
(NH4)2S,  especially  in  the  presence  of  considerable  NH4OH.  Such  a  coloration  is  an 
indication  that  these  reagents  have  been  added  in  unnecessarily  large  amounts.  To 
clear  the  filt.,  acidify  with  HC1,  make  faintly  alkaline  with  NH4OH  and  add,  if  neces- 
sary, a  drop  or  two  of  (NH4)2S  to  bring  down  NiS  and  filter. 

$  Normal  HC1  may  be  made  by  diluting  the  cone,  acid  with  12  volumes  of  water  or 
the  dil.  acid  with  3  volumes. 


THE    BASIC    ANALYSIS.  269 

DISCUSSION  OF  THE  SEPARATION  AND  ANALYSIS  OF  GROUP  III. 

The  separation  of  Co,  Ni,  Mn  and  Zn  as  Group  III  depends  upon  the  facts 
that  the  sulfids  of  these  elements,  although  soluble  in  HC1,  are  insoluble  in  water, 
and  secondly,  that  the  hydroxids  are  not  precipitated  by  NH4OH  in  the  presence 
of  NH4C1,  owing  to  the  repression  of  the  ionization  of  the  weak  base  by  its  more 
readily  dissociating  salts  (see  page  31).  For  these  reasons  Co,  Ni,  Mn  and  Zn 
are  not  precipitated  with  the  sulfids  of  Group  I,  Divs.  B  and  C.  The  sulfids  and 
hydroxids  of  Groups  IV  and  V  either  do  not  form  or  are  soluble  under  those 
conditions. 

The  precipitated  sulfids  are  washed  with  dilute  NH4C1  solutions  instead  of  pure 
water  to  prevent  the  formation  of  pseudo-solutions  toward  which  these  sulfids 
show  a  strong  tendency.  A  few  drops  of  (NH4)2S  are  usually  added  to  the  NH4C1 
to  prevent  oxidation  of  moist  CoS  and  NiS  to  soluble  sulfates.  For  this  reason 
the  precipitate  should  be  washed  and  treated  with  acid  as  quickly  as  possible. 

The  separation  of  Co  and  Ni  from  Mn  and  Zn  depends  on  the  relative  insolu- 
bility of  CoS  and  NiS  in  cold  N  HC1.  Traces  of  these  elements  only  dissolve 
with  MnS  and  ZnS,  not  great  enough  to  interfere  with  tests  for  the  latter  ele- 
ments. As  all  the  sulfids  are  more  soluble  in  warm  HC1,  heat  should  be  avoided 
unless  a  light-colored  precipitate  indicates  that  both  Co  and  Ni  are  absent. 

The  detection  of  Co  is  not  at  all  difficult  even  in  the  presence  of  many  times 
as  much  Ni,  the  borax-bead  test  being  exceedingly  delicate.  The  detection  of  Ni 
in  the  presence  of  Co  is  more  difficult.  Aqua  regia  oxidizes  both  sulfids  to  sul- 
fates, thus  rendering  them  soluble.  The  separation  by  means  of  K3(Fe  =  C6N6) 
depends  upon  the  fact  that  in  the  presence  of  NH4OH,  a  soluble  nickel-ammonia 
ferricyanid  is  formed  whereas  the  cobalt  compound  produced  under  the  same 
conditions  is  insoluble.  The  maroon  precipitate  is  a  cobaltous  ferricyanid  con- 
taining ammonia,  the  proportion  between  the  constituents  apparently  varying 
somewhat.  The  separation  may  be  made  most  satisfactorily  in  solutions  acid 
with  acetic  but  not  with  mineral  acids.  The  aluminum  salt  is  added  simply  that 
the  slimy  Co3(Fe  =  CQN6)2  may  be  entangled  in  the  A1(OH)3  precipitated  by 
NH4OH  and  thus  readily  filtered.  The  addition  of  NaOH  to  the  filtrate  trans- 
poses nickelous  ferricyanid  into  Ni(OH)2  and  Na3(Fe  =  C6N6).  Since  ferri- 
cyanids  in  alkaline  solution  act  as  oxidizing  agents  (see  page  131),  nickelous 
hydroxid  is  oxidized  into  nickelic  hydroxid,  Ni(OH)3. 

Filt.  24  is  boiled  to  remove  H2S  which,  if  present,  would  precipitate  the 
sulfids  again  on  making  the  solution  alkaline.  The  separation  of  Mn  and  Zn 
depends  upon  the  fact  that  of  the  hydroxids  first  precipitated  by  NaOH,  zinc 
hydroxid  can  be  made  to  play  the  part  of  an  acid  and  will  dissolve  in  an  excess 
of  an  alkaline  hydroxid  to  form  soluble  zincates  (see  page  103).  Manganese 
hydroxid  will  not  act  as  an  acid  and  is  insoluble  under  the  same  conditions.  The 
addition  of  HC1  to  the  alkaline  zincate  sets  free  Zn02H2  and  then  causes  it  to 
play  the  part  of  a  base  again,  in  which  condition  it  may  be  precipitated  as  a 
metallic  sulfid. 

Boiling  with  HN03  and  Pb02  oxidizes  Mn(OH)2  from  the  divalent  base-form- 
ing condition  into  the  heptavalent  acid-forming  condition,  the  change  being 
indicated  by  the  pink  color  of  the  solution.  The  test  is  very  delicate  and  but 
little  Mn(OH)2  should  be  used  as  1  gm.  of  Pb02  will  not  be  sufficient  to  oxidize 
a  large  amount  of  the  precipitate. 

A  similar  oxidation  takes  place  on  fusion  with  Na2C03  and  KN03  but  in  the 
presence  of  alkalies  proceeds  only  as  far  as  the  green  hexavalent  manganates. 


270 


QUALITATIVE    ANALYSIS. 


SEPARATION  AND  ANALYSIS  OF  GROUP  IV. 

Heat  to  boiling  Filt.  23  (which  should  contain  NH4C1  and  be  alkaline  with 
NH4OH*),  add  5-10  ccm.  (NH4)2C03,f  boil  for  a  moment,  add  a  few  drops  of 
(NH4)2C03  again  and  filter  immediately. 

PREC.  27.     Group  IV.     BaCO3,  SrC03,  CaC03.     Treat          FILT.  27.     Group 
on  the  paper  with  as  little  H(C2H302)  as  will  suffice  to       V.     See  page  272. 


dissolve  the  prec.  and  wash  through  the  filter  into  a  test-tube.    Add  a  few  drops 
of  K2Cr207  and  filter. 


PREC.  28.  Yellow 
prec.,  BaCrO4, 
indicates  Ba. 

To  confirm,  dis- 
solve in  HC1. 

A.  Heat  on 
platinum  wire. 
Green  flame  con- 
firms Ba. 

B.  Add  an 
equal  volume  of 
saturated  CaS04 
solution. 

White  prec., 
BaS04  confirms 
Ba. 


FILT.  28.  Make  alkaline  with  NH4OH,  heat  to  boiling, 
add  (NH4)2C03,  filter  and  wash  until  yellow  color  has  en- 
tirely disappeared. 

PREC.  29.     SrC03,  CaC03.    Dissolve  in    I       FILT.  29. 
as  little  HN03  as  possible  and  evaporate       Reject, 
nearly  to  dry  ness.     Add  5-10  ccm.  amyl  alcohol  and  dehy- 
drate by  boiling  until  the  vapor  of  the  alcohol  will  burn 
quietly  at  the  mouth  of  the  test-tube,  taking  care  that  the 
volume  does  not  become  less  than  5  ccm. 

//  no  res.  is  left,  Sr  is  absent.  Prec.  29  must  have  been 
CaCO3.  Confirm  by  testing  as  directed  under  Filt.  30. 

//  a  res.  is  present,  it  indicates  Sr,  or  possibly  traces  of 
Ca.  Carefully  decant  the  liquid  from  the  res.  and  pour 
through  a  dry  filter. 


RES.  30.  Ignite  the  res.  in 
the  test-tube  gently,  cool,  dis- 
solve in  a  drop  or  two  of  HNO,, 
treat  with  5-10  ccm.  of  amyl 
alcohol  and  dehydrate  as  be- 
fore. 

Light  colored  res.  indicates 
Sr.    Moisten  with  a  drop  of  HC1 
and  heat  on  platinum  wire. 

Brilliant  crimson  flame  con- 
firms Sr. 


FILT.  30.  Dilute 
with  an  equal  volume 
of  ordinary  ethyl  al- 
cohol and  add  H2S04. 

White,floccy  prec.,^: 
CaS04,  indicates  Ca. 

To  confirm,  filter, 
moisten  prec.  with 
HC1  and  heat  on 
platinum  wire. 

Orange  flame  con- 
firms Ca. 


*  In  performing  a  partial  analysis  of  Group  IV  alone  or  with  Group  V,  3-5  ccm. 
of  NH4C1  should  be  added  and  the  solution  made  alkaline  with  NH4OH  before  treating 
with  (NH4)2C03. 

f  Commercial  (NH4)2CO3  consists  mainly  of  a  mixture  of  acid  ammonium  carbon- 
ate and  ammonium  carbamate,  both  of  which  form  soluble  salts  with  the  metals  of  this 
group.  On  heating  to  60°,  the  carbamate  is  changed  into  the  carbonate  according  to 
the  equation  : 

H4N  —  O  H4N  —  O 


H2N 


H4N  —  O 


The  NH4OH  added  to  the  solution  before  precipitation  changes  the  acid  carbonate 
into  normal  carbonate  and  makes  the  solution  alkaline.  Were  (NH4)2CO3  added  to  an 
acid  solution,  CO2  would  be  liberated  which  would  tend  to  form  acid  carbonates,  also. 

$  A  faint  cloudiness  may  be  due  to  Sr. 


THE    BASIC    ANALYSIS.  271 

DISCUSSION  OF  THE  SEPARATION  AND  ANALYSIS  OF  GROUP  IV. 

The  separation  of  Group  IV  depends  on  the  relative  insolubility  of  the  carbon- 
ates-of  Ba,  Sr  and  Ca,  even  in  the  presence  of  NH4C1.  Under  these  conditions 
Mg  forms  a  complex  salt,  (NH4)2MgCl4,  which  does  not  separate  Mg  ions  in 
sufficient  concentration  to  precipitate  the  carbonate  when  treated  with  (NH4)2C08. 
The  carbonates  of  the  alkalies  arc  all  readily  soluble  and  with  Mg  pass  through 
into  the  nitrate. 

Precipitation  from  boiling  solution  and  subsequent  short  boiling  changes  the 
amorphous  precipitate  into  the  crystalline  form  in  which  condition  it  is  much 
more  readily  filtered.  Boiling  should  not  be  continued  longer  than  is  neces- 
sary to  effect  the  change  as,  otherwise,  a  reversion  of  the  reaction  by  which  "the 
carbonates  were  precipitated  takes  place,  owing  to  the  decomposition  of 
(NH4)2C03  into  NH3,  C02  and  H20  which  are  boiled  out  of  solution  (see  page  76). 
In  this  way  a  considerable  amount  of  the  precipitate  may  dissolve  making  the 
separation  of  Group  IV  far  from  exact.  The  following  equations  serve  to  illus- 
trate this  point : 

CaCl,  +  (NH4)0C03  ->  CaC03  +  2NH4C1 
CaC03  +  2NH4C"l  ->  CaCl2  +  (NHJ2C03  -»  2NHZ+  C02  +  H20 

The  separation  of  Ba  from  Sr  and  Ca  depends  on  the  fact  that  a  solution  of 
K2Cr207  contains  sufficient  Cr04  ions  (see  page  195)  to  precipitate  BaCr04  but  not 
SrCr04  or  CaCr04.  In  alkaline  solution,  the  ions  Cr207  change  into  Cr04  ions 
which  may  precipitate  SrCr04  from  concentrated  solutions,  hence  the  precipi- 
tation is  made  in  the  presence  of  acetic  acid  in  which  BaCr04  is  not  appreciably 
soluble. 

Before  separating  Sr  from  Ca,  it  is  necessary  to  remove  chromates  entirely; 
for  these  salts  are  insoluble  in  amyl  alcohol  and  will  remain  with  the  Sr.  Ca  and 
Sr  are  therefore  reprecipitated  as  carbonates  and  washed  until  the  yellow  color 
has  entirely  disappeared.  Their  separation  depends  on  the  fact  that  anhydrous 
Sr(N03)2  is  insoluble  in  boiling  amyl  alcohol  while  Ca(N03),  is  soluble.  Dehy- 
dration of  the  nitrates  is  readily  accomplished  by  boiling  with  amyl  alcohol 
because  its  boiling-point  (128-132°)  is  higher  than  that  of  water.  As  long  as 
water  is  present,  the  alcohol  will  boil  with  slight  explosions  and  the  vapor  will 
not  burn  quietly. 

By  the  action  of  free  HN03  (a  little  of  which  is  always  present  after  dis- 
solving the  carbonates)  on  amyl  alcohol,  certain  organic  acids  may  be  formed 
capable  of  uniting  with  Ca  and  preventing  solution.  Any  basic  nitrate  of  Ca, 
which  may  sometimes  form,  is  likewise  insoluble.  Hence  it  is  necessary  to  repeat 
the  extraction  with  amyl  alcohol  to  be  certain  that  only  Sr  remains  undissolved. 
Ignition  of  the  residue  is  necessary  to  decompose  the  salt  of  any  organic  acid 
into  the  carbonate  which  is  then  converted  into  nitrate  by  HN03. 

Traces  of  Sr  may  dissolve  in  amyl  alcohol  and  give  a  slight  cloudiness  to  the 
filtrate.  This  will  scarcely  interfere  with  the  test  for  Ca  inasmuch  as  CaS04 
comes  down  in  a  heavy  flocculent  precipitate  under  these  conditions.  Amyl 
alcohol  and  water  are  not  miscible.  Ethyl  alcohol  mixes  readily  with  either.  By 
adding  an  equal  volume  of  the  latter  solvent  a  homogeneous  mixture  may  be 
obtained  when  dilute  H2S04  is  added. 


272 


QUALITATIVE    ANALYSIS. 


ANALYSIS  OF  GROUP  V. 

To  Filt.  27  add  a  few  drops  of  (NH4)2S04  and  (NH4)2C204,  warm  and  filter  if 
necessary. 

PREC.  31.*  BaS04,  CaC2O4,  (SrS04).   I        FILT  31.     Divide  into  two  portions. 


A.  AddHNa2P04.  White 
prec.,  Mg(NH4)P04,  indi- 
cates Mg. 

Confirm  by  dissolving  the 
prec.  suspended  in  the  liquid 
in  the  least  possible  amount 
of  HC1,  heat  to  boiling  and 
neutralize  carefully  by  the 
addition  of  NH4OH,  drop 
by  drop.  The  reappearance 
of  the  prec.  in  crystalline 
form  confirms  Mg. 


B.  Evaporate  to  dryness  and  ignite  until  am- 
monium salts  are  entirely  removed,  f  i.  e.}  white 
fumes  cease  to  be  given  off. 

(1)  //  Mg.  was  not  found  in  A,  test  for  the 
alkali  metals  as  directed  under  Filt.  33. 

(2)  //  Mg  was  found  in  A,  boil  out  the  ignited 
res.  with  a  little  water,  add  Ba(OH)2  until  the 
solution  is  alkaline  to  litmus  and  filter. 


PREC.  32. 
Mg(OH)2.    Reject. 


PREC.  33. 
BaS04.     Reject. 


FILT.  32.  Add  H2S04 
until  the  solution  is  just  acid 
to  litmus,  warm  and  filter. 

FILT.  33.  Evaporate  al- 
most to  dryness  and  test  for 
K,  Na  and  Li  by  heating  on 
a  platinum  wire  in  a  Bunsen 
flame. 


Examined  with  an  ordinary  spectroscope  Na,  if  present  in  solution,  gives  a 
single  yellow  line  of  great  intensity.:}:  Li  gives  a  crimson  line  a  short  distance  to 
the  left  of  the  Na  line.  K  gives  a  red  line  still  further  to  the  left  of  the  Na  line 
than  is  the  Li  line  and  a  blue  line  far  toward  the  right  end  of  the  spectrum.  See 
Table  of  Spectra,  frontispiece. 

Examined  without  the  aid  of  the  spectroscope  Na  gives  an  intensely  yellow 
flame,  Li  a  carmine  flame  and  K  a  violet  flame. § 

*  If  this  prec.  is  appreciable  and  Prec.  27  was  unsatisfactory  in  quantity  owing 
to  the  solvent  action  of  NH4C1  on  the  carbonates,  Prec.  31  may  be  rendered  soluble  by 
fusion  with  Na2CO3  as  directed  on  page  262.  The  melt,  after  extraction  with  boiling 
water,  may  be  treated  as  Prec.  27,  and  examined  for  members  of  Group  IV. 

t  Ammonium  salts  are  introduced  during  the  course  of  the  analysis,  hence  must  be 
tested  for  before  beginning  the  regular  analysis  (see  page  249).  If  Na  or  K  have 
been  introduced  in  the  preparation  of  the  solution  (see  page  241)  or  otherwise,  their 
detection  in  Group  V  is  without  significance.  Special  treatment  must  be  given  in 
order  to  ascertain  their  presence.  The  original  substance  may  be  extracted  with  boiling 
water  or  HC1  and  the  resulting  solution  tested,  or,  in  the  case  of  insoluble  silicates,  the 
method  given  on  page  242  may  be  followed. 

J  Because  of  the  extreme  delicacy  of  this  method  of  detection,  the  Na  line  is  ever- 
present  with  greater  or  less  intensity  even  when  Na  is  not  present  in  solution  in  appre- 
ciable quantity. 

§  In  the  presence  of  Na,  the  K  flame  is  masked  and  cannot  be  detected  with  the 
unaided  eye.  Viewed  through  a  potassioscope  (see  page  57),  however,  it  is  readily 
distinguishable. 


THE    BASIC   ANALYSIS.  273 

DISCUSSION   OF   ANALYSIS    OF   GROUP   V. 

Group  V  consists  of  those  metals  that  are  not  precipitated  as  chlorids,  sulfids, 
basic  acetates  or  carbonates  under  the  conditions  of  the  analysis  and  are  therefore 
left  together  as  a  soluble  residue  after  each  of  the  preceding  groups  has  been 
precipitated. 

Magnesium  differs  in  its  chemical  nature  from  the  other  members  of  this 
group  and  can  be  separated  from  them  by  precipitation  as  the  phosphate,  for 
phosphates  of  the  alkali  metals  only  are  soluble.  This  reaction  for  the  detection 
of  Mg  is  of  value  only  in  the  absence  of  other  metals,  for  every  one  of  the  metals 
considered  in  the  present  scheme  of  analysis  will  form  a  precipitate  when  treated 
with  HNa2P04  in  ammonical  solution  with  the  exceptions  of  As  and  Sb.  Of  these 
the  metals  of  Group  IV  only  are  likely  to  be  present  in  appreciable  quantities, 
owing  to  the  solvent  action  of  NH4C1  on  the  carbonates  of  Ba  and  Ca  especially. 
The  sulfate  and  oxalate  ions  remove  these  metals  from  solution  so  completely, 
however,  that  there  is  no  possibility  of  precipitating  them  as  phosphates,  for 
BaS04  and  CaC204  are  much  less  soluble  than  the  corresponding  phosphates. 
Strontium  as  sulfate  or  oxalate  is  not  likely  to  be  present  unless  Prec.  31  be  very 
considerable  in  quantity  thus  indicating  that  Prec.  27  was  not  brought  down  under 
satisfactory  conditions.  Strontium  carbonate  is  much  less  soluble  in  water  than 
either  BaC03  or  CaC03  and,  therefore,  is  much  less  affected  by  NH4C1. 

The  absence  of  all  other^  metals  being  thus  assured,  any  precipitate  which 
falls  after  introducing  the  P04  ions  must  be  due  to  Mg.  Small  amounts  of  this 
phosphate  are  slow  in  separating  from  supersaturated  solutions,  however,  but,  on 
standing  over  night,  a  crystalline  precipitate  will  settle.  If  the  inside  of  the  glass 
vessel  holding  the  solution  be  scratched  with  a  glass  rod,  the  crystals  tend  to  collect 
along  the  roughened  surface  and  may  be  more  readily  detected. 

The  exceedingly  close  chemical  similarity  between  K,  Na  and  Li  renders  any 
separation  of  these  elements  difficult.  Fortunately,  however,  they  may  be  readily 
detected  by  their  characteristic  flame  colors.  In  the  spectroscopic  examination, 
both  Mg  and  NH4  salts  tend  to  make  the  indications  less  distinct  and  they  are  re- 
moved therefore.  All  ammonium  salts  are  volatile  and  may  be  driven  off  by 
ignition.  After  the  removal  of  NH4C1,  Mg  may  be  precipitated  by  the  OH  ion, 
for  the  introduction  of  which,  obviously,  neither  NaOH,  KOH  or  NH4OH  will 
serve.  Barium  hydroxid,  Ba(OH)2,  is  used  because  any  excess  may  be  readily 
and  completely  removed  by  the  addition  of  H2S04.  In  the  resulting  filtrate,  after 
concentration,  the  alkali  metals  may  be  readily  detected  spectroscopically.  Inas- 
much as  the  spectroscope  reveals  the  presence  of  exceedingly  minute  quantities  of 
these  elements  and  since  Na  salts  are  present  everywhere  as  a  fine  dust  in  the  air,  it 
is  necessary  to  judge  carefully  the  intensity  of  the  indication  in  deciding  whether 
or  not  Na  salts  are  present  in  solution  in  appreciable  quantities.  This  applies  to 
the  examination  of  the  flame  color  with  the  unaided  eye  as  well.  If  Na  be  present, 
the  flame  will  be  intensely  yellow  and  the  indication  will  be  lasting. 


19 


274  QUALITATIVE   ANALYSIS. 


THE  ACID  ANALYSIS.  275 


THE   ACID   ANALYSIS. 

Owing  to  the  fact  that  anions  show  a  much  greater  variety  of  reactions  than 
do  kathions,  the  Acid  Analysis  cannot  be  made  as  systematic  as  the  Basic  Analysis. 
The  method  of  the  Basic  Analysis  cannot  be  applied  to  the  acids.  The  separation 
of  the  acids  by  precipitation  into  groups  from  which,  by  further  analysis,  each 
individual  acid  may  be  isolated  and  detected  can  not  be  accomplished  satisfactorily. 
The  method  of  the  Acid  Analysis  is,  in  general,  to  test  the  prepared  solution  for 
the  presence  of  each  acid  by  individual  tests.  By  certain  preliminary  reactions, 
however,  the  presence  or  absence  of  certain  groups  may  be  indicated.  The  classi- 
fication is  dependent  mainly  on  the  action  of  the  Ba  and  Ag  ions,  supplemented 
by  the  action  of  certain  reducing  and  oxidizing  agents.  The  groups  thus  formed 
are  not  all  mutually  exclusive  but  certain  acids  may  appear  in  two  or  more  groups. 
Careful  consideration  of  the  implications  of  these  preliminary  tests  and  intelli- 
gent forethought  will  simplify  and  shorten  the  Acid  Analysis  very  considerably. 


Acids  precipitated  as  insoluble  salts  by  the  addition  of  the  Ba  ion  to  a  neutral 
solution : 

H2S04,  H2SiF6,  HF,  H2C204,  H2Si03,  H2Cr207,  H2CO3,  H3P04,  H3As04, 
H3As03,  H3B03,  H2S03,  H2S2O3,  H2(C4H406). 

Acids  precipitated  as  insoluble  salts  by  the  addition  of  the  Ag  ion  to  solutions 
acidified  with  HN03 : 

H3(Fe  «  C6N6),  H4(Fe  =  C6N6),  HSCN,  HCN,  H2S,  HI,  HBr,  HC1, 
HOC1,  HN02,*  H(C2H302).* 

Acids  not  precipitated  under  either  of  the  above  conditions,  the  barium  salts 
being  soluble  in  water  and  the  silver  salts  in  dilute  nitric  acid : 

H(C2H302),*  HN02,*  HN03,  HC1O3,  HC104,  HMn04,  HOCN,  H202. 


*  The  silver  salts  of  HNO2  and  H(C2H3O2)  can  be  precipitated  from  concentrated 
solutions  only. 


276 


QUALITATIVE    ANALYSIS. 


PRELIMINARY    TESTS    FOR    THE    PRESENCE 


(1)  Acidify  a  portion  of  the  prepared  solution  with  HC1.     A  prec 
sulfur,  or  (b)  iodin  indicates  the  presence  of  one  or  more  of  A  or  B. 


OF    ACIDS. 

of  (a) 


A. 
H2S203. 
H2SX. 
H2S  + 

Oxidizing 
agents. 

Filter  clea 
C. 
H2S04. 
H2SiFa. 

r  and  add  BaC 
Filter  cle 
CaCl2  and  j 
indicates  on 
D. 
HF. 
H,C,04. 
H,Si03. 
H2Cr207 
if  prec.  is 
yellowish. 

2.  White  prec.  indicates  one  or  more  of  C. 
ar  and  boil  to  remove  C02.     Cool,  add 
m  equal  volume  of  Na(C2H302).    Prec. 
e  or  more  of  D. 
Filter  clear  and  make  alkaline  with 
Ba(OH)2.      Prec.    indicates    one    or 
more  of  E. 

E. 
H2COS.         H3As03.       H,S03. 
H3P04.         H0Si03.        H0S,03. 
H3As04.       H3B03.         H2(C4H406). 

HI  + 

Oxidizing 
agents. 

(2)  Acidify  a  second  portion  of  the  prepared  solution  with  HN03  and  add 
AgN03.  A  prec.  indicates  one  or  more  of  F  and  possibly  one  or  more  of  G. 

Absence  of  a  prec.  indicates  the  absence  of  all  of  F  but  not  necessarily  of 
all  of  G. 


F.— 

H3(Fe-C6N6)- 

H4(Fe  =  C6N6). 

HSCN. 

HON. 

H2S. 


HI. 
HBr. 
HC1. 
HOC1. 


G.— 
HN02. 
H(C2H3OJ. 


H.— 

L— 

H4(Fe  =  CfiNR). 

4  x                         6       6' 

HN03. 

H,S. 

HC103 

HI. 

HC104. 

HBr. 

HMn04. 

HOCN. 

H202. 

If  a  prec.  was  obtained  with  AgN03,  shake  thoroughly  to  coagulate  it  and  allow 
to  settle.  Pour  off  the  liquid  and  wash  the  prec.  by  decantation.  Treat  the  prec. 
with  dil.  NH4OH  and  shake  thoroughly.  If  the  prec.  is  entirely  dissolved,  no 
member  of  H  can  be  present. 

(3)  No  prec.  with  AgN03  nor  with  BaCl2  indicates  the  absence  of  all  A,  B, 
C,  D,  E  and  F.    One  or  more  of  G  or  I  may  be  present,  however. 

(4)  Test  for  oxidizing  acids:    To  1-2  ccm.  of  the  prepared  solution  add 
twice  its  volume  of  MnCl2  in  cone.  HC1  and  warm  gently.     The  appearance  of  a 
dark  coloration  (higher  chlorids  of  manganese)  is  due  to  the  presence  of  some 
oxidizing  agent.     One  or  more  of  J  may  be  present.     The  absence  of  such  a  color 
proves  the  absence  of  all  of  J. 

J.  H202,  HN02,  HN03,  HOC1,  HC103,  HMn04,  H2Cr2O7,  H3(Fe^C6N6). 

(5)  Test  for  reducing  acids:    Acidify  a  portion  of  the  prepared  solution 
with  H2S04  and  add  a  single  drop  of  dil.  KMn04  solution.     If  color  is  not  bleached 
after  ten  minutes,  none  of  K  is  present. 


K.— 
H2S. 
H,S03. 
H  S003. 

HSCN. 


HI. 
HBr. 


HN0. 


=  C6N6).H202. 


If  color  is  not  bleached  in  the  cold  solution,  heat 
nearly  to  boiling.  If  the  color  still  remains  none 
of  L  is  present. 


L.— 


H,(C4H406). 


If  bleaching  does  not  occur  either 
hot  or  cold,  one  or  more  of  M  may  be 
present  but  none  of  K  or  L. 


M.— H2S04>  H0SiF6,  HF,  H  Si03,  H2Cr0O7,  H2C03,  H3P04,  H3As04,  H3BO3, 

H3(Fe=C6N6),  HCN,  HC1,  HOC1,  H(C,H3O2),  HNO3, 

HC103,  HC104,  HMn04,  HOCN. 


THE  ACID  ANALYSIS.  277 

DISCUSSION    OF    THE    PRELIMINARY    TESTS    FOR   ACIDS. 

While  the  salts  of  certain  acids  are  stable  in  neutral  or  alkaline  solution,  the 
free  acids  are  unstable  and  give  evidence  indicative  of  their  presence.  When  the 
solution  is  acidified  sulfur  is  separated  (on  standing)  if  thiosulfates  or  polysulfids 
are  present,  or  sulfids  together  with  oxidizing  agents.  Similarly  iodids  with 
oxidizing  agents  separate  iodin.  These  acids  constitute  groups  A  and  B. 

Group  C  consists  of  those  acids  whose  barium  salts  are  insoluble  in  HC1,  and 
group  D  of  those  whose  barium  or  calcium  salts  are  insoluble  in  acetic  acid.  By 
the  addition  of  sodium  acetate,  the  solution  acid  with  HC1  is  changed  to  one  acid 
with  H(C2H302).  Calcium  chlorid  is  added  inasmuch  as  the  calcium  salts  of  HF 
and  H2C204  are  less  soluble  than  the  corresponding  barium  salts.  Silicic  acid 
may  or  may  not  come  down  in  the  presence  of  acetic  acid,  its  action  depending 
largely  on  the  amount  present, 

Group  E  consists  of  those  acids  whose  barium  or  calcium  salts  are  precipitated 
from  alkaline  solutions,  Ba(OH)2  being  used  in  preference  to  NH4OH  since  the 
salts  of  the  last  four  members  of  this  group  are  appreciably  soluble  in  the  presence 
of  ammonium  salts.  Sodium  or  potassium  hydroxids  always  contain  carbonates. 

Group  F  contains  the  acids  whose  silver  salts  are  highly  insoluble  in  HN03. 
Hypochlorous  acid  appears  in  this  group  because  its  silver  salt,  although  soluble 
in  water,  decomposes  very  readily,  according  to  the  equation : — 

SAgOCl  -»  2AgCl  -f  AgCIO,. 

Group  G  consists  of  those  acids  whose  silver  salts  under  the  same  conditions 
may  be  precipitated  only  from  concentrated  solutions.  The  acids  of  these  two 
groups  may  be  classified  in  another  way,  owing  to  the  fact  that  the  silver  salts  of 
certain  of  them  are  not  readily  soluble  in  NH4OH.  These  are  placed  together, 
therefore,  in  group  H.  Silver  bromid  is  appreciably,  but  usually  not  entirely, 
soluble  in  NH4OH. 

Group  I  consists  of  those  acids  whose  barium  salts  are  soluble  in  water  and 
whose  silver  salts  are  not  precipitated  in  the  presence  of  HN03. 

Individual  tests  must  be  made  for  each  of  the  acids  in  the  groups  indicated 
to  be  present  by  these  preliminary  experiments  unless  the  tests  for  oxidizing  and 
reducing  acids  indicate  that  certain  of  them  are  not  present. 

Under  the  influence  of  oxidizing  agents  in  general,  HC1  is  oxidized  to  water 
and  free  chlorin.  In  concentrated  HC1  solution,  the  higher  chlorids  of  manganese, 
MnCl3  or  MnCl4  (see  page  214),  dark  brown  in  color,  may  be  formed  and  are 
capable  of  temporary  existence.  The  acids  which  act  as  oxidizing  agents  under 
these  conditions  constitute  group  J.  Ferricyanic  acid  gives  a  dark-drab,  opaque 
solution  appearing  somewhat  different  from  the  other  acids. 

Many  free  acids  act  as  reducing  agents  and  will  bleach  a  dilute  solution  of 
KMn04.  Such  as  act  in  the  cold  are  classed  together  in  K  while  oxalic  and  tartaric 
acids  usually  require  heating  and  constitute  group  L.  Hydrocyanic  acid  may 
cause  slight  bleaching  in  hot,  fairly  concentrated  solutions  but  most  of  the  HCN 
will  boil  out  unchanged  and  may  be  detected  by  its  characteristic  odor.  Most  acids 
do  not  have  any  reducing  action  and  are  classed  as  group  M.  Two  bodies,  H202 
and  HN02,  because  of  their  instability  react  readily  with  either  oxidizing  or  reduc- 
ing agents  and  therefore  are  found  both  in  groups  J  and  K.  Although  salts  are 
much  more  stable  than  free  acids,  in  general  oxidizing  and  reducing  agents  can- 
not exist  in  the  same  solution. 

If  no  test  for  any  acid  can  be  obtained,  the  original  substance  must  have  been 
an  element  or  its  oxid  or  hydroxid. 


278  QUALITATIVE    ANALYSIS. 

TEST  FOR   SULFUEIC  ACID. 

Acidify  a  portion  of  the  prepared  solution  with  HC1,  add  BaCl2  as  long  as  a 
precipitate  falls,  filter,  wash  and  dry  the  precipitate.*  Mix  with  Na2C08  and 
heat  before  the  blowpipe  upon  charcoal.  Remove  the  residue  from  the  charcoal, 
place  upon  a  silver  coin  and  moisten  with  a  drop  of  water. 

A  dark  stain,  Ag2S,  indicates  the  presence  of  H2S04. 

*  Part  of  this  precipitate  may  be  used  in  testing  for  H2SiF6  (see  next  test). 


TEST    FOR  HYDROSILICOFLUORIC    ACID. 

Acidify  a  portion  of  the  prepared  solution  with  HC1,  add  BaCl2  as  long  as  a 
precipitate  falls,  stand  for  ten  minutes  then  filter,  wash  and  dry  the  precipitate 
thoroughly.  Heat  in  a  test-tube  with  cone.  H2S04  and  hold  in  the  fumes  evolved 
a  loop  of  platinum  wire  containing  a  drop  of  water. 

The  formation  of  a  white  deposit,  H.,SiO3,  in  the  water  indicates  the  presence 
of  H2SiT6. 


TESTS  FOR  HYDROFLUORIC  ACID. 

(1)  In  a  platinum  crucible  treat  a  small  portion  of  the  dry  substance  with 
warm  cone.  H2S04  and  cover  with  a  watch-glass  which  has  been   previously 
coated  with  paraffine  in  which  two  or  three  scratches  have  been  made  with  a 
knife  blade. 

On  standing,  etching  of  the  glass  where  exposed  to  the  fumes  indicates  HF. 

(2)  From  a  portion  of  the  prepared  solution  acidified  with  HC1  remove 
H2S04  and  H2SiF6  by  treating  with  BaCl2  and  filtering.    Warm  the  clear  filtrate 
and  add  CaCl2.    Make  the  solution  alkaline  with  Ba(OH)2  and  after  a  moment 
make  just  acid  again  with  HC1. 

A  white  precipitate,  CaF2,  insoluble  in  dilute  HC1,  indicates  HF. 

Confirm  by  filtering  and  mixing  the  dried  precipitate  with  an  equal  amount 
of  a  dry  silicate.  Treat  with  cone.  H0S04,  exposing  a  rod  moistened  with  water 
to  the  fumes  evolved. 

White  precipitate,  appearing  as  flecks  in  the  drop  of  water,  H2Si03,  con- 
firms the  test  for  HF. 


THE  ACID  ANALYSIS.  279 

DISCUSSION   OF    THE    TEST    FOR  SULFURIC   ACID. 

In  the  presence  of  HC1  the  barium  ion  gives  no  precipitate  except  with 
sulfuric  acid  and  hydrosilicofluoric  acid  (which  see).  The  formation  of  a 
precipitate  may  indicate  the  presence  of  both  of  these  acids  or  of  one  in  the 
known  absence  of  the  other.  Fusion  with  Na2C03  oh  charcoal  results  in  the 
reduction  of  the  sulfate  to  the  sulfid,  which  when  moistened  and  brought  in 
xjontact  with  metallic  silver  produces  a  dark  stairi  of  silver  sulfid.  Oxygen  is 
absorbed  from  the  air  at  the  same  time  according  to  the  equation: 

4Ag  +  2Na2S  +  02  +  2HaO  ->  2Ag,S  +  4NaOH 


DISCUSSION    OF    THE    TEST    FOR    HYDROSILICOFLUORIC    ACID. 

Under  the  conditions  of  the  experiment,  BaSiF6  (or  BaSOJ  <M  'y  can  be 
precipitated.  When  the  dried  precipitate  is  heated  with  cone.  H^y-J,,  H2SiFa 
is  liberated  and  decomposes  into  SiF4  and  HF,  both  of  which  a  -e  volatile. 
On  coming  in  contact  with  water  SiF4  reacts  with  it  to  form  silicic  acid  which 
appears  as  white  flecks  within  the  liquid  (see  page  133).  The  mix'nn-  in  the 
test-tube  must  be  anhydrous  as  otherwise  this  reaction  will  take  pla«.'>  immedi- 
ately and  SiF4  will  never  be  volatilized. 


DISCUSSION    OF    THE    TESTS    FOR   HYDROFLUORIC  ACID. 

Test  (1)  depends  on  the  corrosive  action  of  HF  on  glass  and  other  sili- 
cates (see  page  199).  It  is  a  delicate  test  and  on  melting  off  the  paraffine,  etching 
of  the  glass  will  appear  along  the  lines  where  the  paraffine  was  scratched. 

The  separation  of  HF,  as  the  second  test,  depends  upon  the  fact  that 
although  barium  fluorid  will  not  come  down  in  the  presence  of  HC1,  once  pre- 
cipitated it  is  not  readily  soluble  in  dilute  HC1.  Calcium  fluorid  is  somewhat 
more  insoluble  in  water  and  acid  than  barium  fluorid.  Since  all  other  acids 
giving  barium  salts  insoluble  in  water  are  distinctly  weaker  than  hydrochloric 
acid,  their  salts  will  be  readily  dissolved  away  from  the  fluorids  by  HC1.  Am- 
monium salts  exert  quite  a  solvent  action  on  fluorids,  however.  By  working 
with  the  prepared  solution  (made  by  boiling  with  sodium  carbonate)  and  using 
barium  hydroxid  as  base,  all  ammonium  salts  are  avoided. 


280  QUALITATIVE    ANALYSIS. 

TESTS    FOR    SILICIC   ACID. 

(1)  In  a  platinum  crucible  intimately  mix  some  of  the  dry  substance  with 
an  equal  weight  of  powdered  CaF2,  moisten  with  cone.  H2S04,  and  heat  under 
the  hood.     Expose  a  film  of  water  on  a  loop  of  platinum  wire  to  the  fumes 
evolved. 

The  formation  of  white  flecks  of  silicic  acid  indicates  the  presence  of  H2Si03. 

(2)  Acidify  a  portion  of  the  prepared  solution  with  HC1,  evaporate  to  dry- 
ness  and  desiccate  at  120°   (see  page  263).     Extract  by  boiling  with  dilute  HC1, 
pour  off  the  solution  from  the  residue,  and  wash  by  decantation.    Heat  some  of 
the  residue  before  the  blowpipe  in  a  sodium  meta-phosphate  bead    (made  by 
heating  microcosmic  salt). 

The  formation  of  an  infusible  "  skeleton,"  SiO2,  not  dissolving  in  the  bead, 
indicates  H2SiO8. 


TEST   FOR   DICHROMIC   ACID.* 

Make  a  portion  of  the  prepared  solution  faintly  acid  with  HN03  and  boil 
to  remove  C02.  Make  just  alkaline  with  NH4OH,  add  CaCl2,  warm  and  stand 
aside  for  10-15  minutes,  shaking  vigorously  from  time  to  time  to  induce  pre- 
cipitation. Filter,  acidify  the  clear  filtrate  slightly  with  HN03  and  add  BaCl2. 
Filter  again  and  to  the  clear  filtrate  add  Pb(C2H302)2,  and  Na(C2H302)  if 
necessary  to  effect  precipitation. 

A  yellow  precipitate,  PbCr04,  indicates  H2Cr2Or 

To  confirm,  filter  and  treat  the  precipitate  on  the  filter  paper  with  warm 
HN03,  allowing  the  solution  to  filter  through  into  a  test-tube.  Cool  and  add  a 
few  grains  of  Na202. 

The  appearance  of  a  transitory  blue  color,  due  to  HCr04,  confirms  H,Cr2O7. 

*  Chromic  acid,  H2CrO4,  is  unknown,  for  as  soon  as  it  is  liberated  from  its  salts, 
it  loses  water  and  passes  into  dichromic  acid  (see  page  195). 


THE  ACID  ANALYSIS.  281 


DISCUSSION    OF    TESTS    FOR    SILICIC    ACID. 

All  silicates  are  decomposed  by  HF  with  the  formation  of  SiF,,  hence 
nothing  will  interfere  with  test  (1)  for  silicic  acid.  Since  silicon  fluorid  is 
decomposed  by  water,  forming  silicic  acid  and  hydrofluoric  acid  (see  page  133) 
the  mixture  in  the  crucible  must  be  kept  anhydrous.  Concentrated  sulfuric  acid 
is  added  for  this  purpose  as  well  as  to  liberate  the  HF. 

The  second  test  can  be  relied  upon  to  detect  silicic  acid  in  solution  under 
all  circumstances  (see  page  135).  If  the  dry  substance  is  similarly  treated,  the 
skeleton  indication  will  appear  in  most  cases  when  silicic  acid  is  present  (see 
page  136) ;  yet  the  indication  is  usually  less  definite,  since  the  basic  constituents 
of  the  original  substances  will  dissolve,  often  rendering  the  bead  turbid. 

Inasmuch  as  silicic  acid  is  usually  detected  during  the  Basic  Analysis  it  will 
generally  be  unnecessary  to  make  tests  for  it  in  the  Acid  Analysis. 


DISCUSSION    OF    THE    TEST    FOR    DICHROMIC    ACID. 

The  presence  of  a  chromate  or  dichromate  is  frequently  indicated  during 
the  Basic  Analysis  by  the  change  in  color  from  yellow  or  red  to  green  when 
the  solution  is  treated  with  H,S.  Chromium,  even  if  present  in  the  acidic 
condition,  is  thus  reduced  to  the  basic  condition  (see  pages  253  and  265),  and  is 
precipitated  with  the  metals  of  Group  II.  If  chromium  is  detected  in  the 
Basic  Analysis,  it  does  not  follow,  however,  that  it  was  originally  present  in 
the  acid  condition,  hence  the  necessity  for  this  test  in  the  Acid  Analysis.  It 
is  unnecessary  to  test  a  colorless  solution  for  the  presence  of  dichromic  acid, 
however. 

While  a  large  amount  of  dichromic  acid  will  make  itself  evident  during  both 
the  Basic  and  Acid  Analyses,  when  present  only  in  small  quantity,  considerable 
care  is  required  for  its  detection.  Certain  acids  must  be  previously  removed. 
Carbonates  are  first  destroyed.  Tartaric  acid  (which  tends  to  hold  up  the 
precipitate  of  PbCrOJ  and  other  acids  are  precipitated  as  calcium  salts  in 
alkaline  solution.  Since  ammonium  salts  exert  a  solvent  action  on  tartrates, 
chromates  and  salts  of  other  acids,  the  amount  of  NH4OH  introduced  must  be 
kept  as  small  as  possible.  Calcium  tartrate  shows  strong  tendency  to  form 
supersaturated  solutions  and  vigorous  shaking  is  necessary  to  bring  down  a 
precipitate.  From  the  filtrate  once  again  acidified  with  HN03,  the  Ba  ion  pre- 
cipitates H2S04.  After  filtering,  the  addition  of  Pb(C2H302)2  introduces  the 
Pb  ion  and  changes  the  solution  from  one  acid,  with  HN03  to  one  acid  with 
H(C2H302),  under  which  condition  PbCr04  is  precipitated.  Sodium  acetate 
affects  the  same  change  and  represses  the  ionization  of  the  acetic  acid  formed  by 
the  reaction.  (See  pages  27  and  31.) 

If  iodids  be  present,  PbI2  may  be  precipitated  also  as  golden  spangles.  They 
are  readily  soluble  in  boiling  water,  however,  and  may  thus  be  separated  from 
the  chromate.  The  confirmatory  test  is  discussed  on  page  267. 


282  QUALITATIVE    ANALYSIS. 

TESTS   FOR  BORIC   ACID. 

(1)  A.  Moisten  some  of  the  dry  substance  with  cone.  H2S04  and  heat  before 
the  blowpipe. 

A  characteristic  green  color  indicates  H3BO3  (in  the  absence  of  Ba  and  Cu 
salts). 

B.  If  a  green  color  does  not  appear  in  A,  mix  the  finely  powdered  sub- 
stance with  CaF2  and  KHS04,  moisten  with  a  drop  of  watfer  and  heat  in  a  loop 
of  platinum  wire,  holding  the  head  at  the  edge  and  near  the  bottom  of  the 
Bunsen  flame. 

A  green  color  indicates  HbB03  (in  the  absence  of  Ba  and  Cu  salts). 

(2)  If  the   substance   exists   in   solution,   evaporate   a   portion   to   dryness, 
add  3-4  ccm.  of  methyl  alcohol  and  an  equal  volume  of  cone.  H2S04,  and  warm, 
igniting  the  vapor  (if  necessary  in  the  apparatus  described  on  page  114). 

A  flame,  green  or  edged  with  green  when  traces  only  are  present,  indicates 
H3B03. 

(3)  Acidify  a  portion  of  the  original  or  the  prepared  solution  with  HC1. 
Absorb  in  turmeric  paper  and  dry  at  100°   (by  placing  on  the  side  of  a  beaker 
in  which  water  is  boiling),  repeating  this  treatment  several  times  if  necessary. 

The  appearance  of  a  pink  coloration  turning  dark  green  on  treating  with 
alkalies  indicates  the  presence  of  H3B03. 


TEST   FOR   CARBONIC   ACID. 

Treat  some  of  the  dry  substance,  or  its  solution  in  water,  with  H2S04  and  pass 
the  gas  evolved  through  "lime  water,"  Ca(OH)2. 

A  white  precipitate,  CaCO3,  indicates  the  presence  of  H2CO3. 

If  the  indication  is  not  obtained  at  first,  warm,  and,  if  necessary,  add  more 
acid. 


THE  ACID  ANALYSIS.  283 


DISCUSSION    OF    TESTS    FOR   BORIC    ACID. 

In  the  flame  test  without  the  use  of  alcohol,  it  is  necessary  to  disti ugui •*. 
carefully  between  barium  or  copper  salts  and  compounds  of  boric  acid,  all  of 
which  color  the  flame  green.  Neither  barium  nor  copper  compounds  are  voi» 
tile  under  the  conditions  prescribed  in  Test  (2),  and  therefore  would  n.  u 
interfere  in  this  way.  While  simple  borates  respond  to  Test  (1)A  very 
satisfactorily,  some  silicates  containing  boron  are  not  sufficiently  decomposed 
by  H2S04  to  give  this  test  for  boric  acid,  but  require  the  stronger  action  of  HF 
formed  from  the  interaction  of  acid  potassium  sulfate  and  calcium  fluorid. 
Under  these  conditions  volatile  boron  fluorid,  BF8,  is  formed  which  also  colors 
the  flame  green. 

For  the  discussion  of  Test  (2)  see  page  114. 

Oxidizing  agents  (except  HN03)  interfere  with  the  turmeric-paper  test 
for  boric  acid,  by  destroying  the  turmeric.  Since  hydriodic  acid  is  readily  de- 
composed by  the  oxygen  of  the  atmosphere  (see  page  211)  with  the  liberation 
of  iodin  which  is  a  bleaching  agent,  the  presence  of  iodids  also  interferes  with 
this  test.  Iodids  may  be  removed  by  treating  the  acid  solution  with  potassium 
nitrite  and  boiling  (see  test  for  iodids,  page  296),  after  which  any  oxidizing 
agents  may  be  reduced  in  acid  solution  by  treatment  with  sodium  sulfite. 
(See  page  185.) 


DISCUSSION    OF    TEST    FOR    CARBONIC    ACID. 

If  salts  of  carbonic  acid  be  present  in  any  appreciable  amount,  the  fact 
will  usually  be  indicated  during  the  Basic  Analysis  as  soon  as  the  solution  is 
acidified.  Therefore,  it  will  be  unnecessary,  generally,  to  make  a  direct  test  for 
carbonic  acid  during  the  Acid  Analysis. 

Some  carbonates  of  metals  forming  insoluble  sulfates  are  not  readily  decom- 
posed by  H2S04,  for  the  reason  that  the  insoluble  precipitate  formed  by  H2S04 
covers  over  the  dry  substance  and  protects  it  from  further  action  by  the  acid. 
If  the  dry  substance  does  not  dissolve  and  no  test  for  carbonic  acid  is  obtained, 
it  may  be  desirable  to  substitute  HC1  and  repeat  the  test. 

Since  an  excess  of  C02  is  liable  to  form  the  soluble  acid  calcium  carbonate 
(see  page  95),  the  formation  of  a  precipitate  should  be  looked  for  shortly  after 
the  bubbles  begin  to  pass  through  the  "  lime  water,"  as  otherwise  a  precipitate 
might  be  formed  and  dissolved  again  without  being  noted. 

In  the  presence  of  sulfurous  acid,  carbonic  acid  may  escape  detection  since 
the  volatile  anhydrid,  S02,  will  pass  off  with  the  C02.  If  S02  is  present  in 
large  quantity,  it  may  neutralize  the  lime  water  and  even  make  the  solution  acid 
before  CO2  sufficient  to  produce  a  precipitate  is  liberated.  If^the^oj^or..of  S02 
is  apparent  while  making  this  test  and  no  indication  of  carbonic  acicl  is  noted, 
it  will  be  safest  to  repeat  the  test  on  a  fresh  portion  of  the  solution.  Iodin 
water  is  added  to  the  acidified  solution  until  the  color  becomes  permanent,  thus 
indicating  that  all  H2S03  has  been  oxidized  to  H2SO4.  The  excess  of  iodin  is 
bleached  with  a  drop  of  SnCl2  and  the  solution  is  then  heated  to  drive  the  CO, 
into  the  "  lime  water." 


284  QUALITATIVE    ANALYSIS. 

TEST    FOR    OXALIC    ACID. 

Faintly  acidify  a  portion  of  the  prepared  solution  with  acetic  acid,  warm  and 
add  an  equal  volume  of  a  saturated  solution  of  CaS04. 

A  white  precipitate,  CaC2O4,  indicates  the  presence  of  H2C2O4. 

To  confirm  filter  and  wash  the  precipitate,  dissolve  in  warm  dilute  H2S04 
and  add  a  drop  of  a  dilute  solution  of  KMn04. 

Bleaching  of  the  permanganate  confirms  the  test  for  H2C204. 


TESTS   FOR   TARTARIC   ACID. 

(1)  Acidify  a  small  portion  of  the  aqueous  solution  with  HC1  and  evapo- 
rate to  dryness.     (If  the  original  substance  is  not  soluble  in  water,  use  the  HC1 
extract.)     Ignite  the  residue  thus  obtained,  noting  carbonization  or  its  absence. 

Charring  denotes  the  presence  of  non-volatile  organic  matter  which  may  be 
tartaric  acid.     Absence  of  carbonization  indicates  the  absence  of  H2(C4H406). 

(2)  If  the  original  substance  contains  carbonizing  matter,  to  10   ccm.   of 
the  prepared  solution  add  a  few  granules  of  dry  K2C03  or  sufficient  to  make  the 
solution  alkaline   (Na2C03  will  not  do).     Evaporate  until  solid  particles  begin 
to  separate.     Dissolve  the  crystals  in  as  little  water  as  possible  and  acidify 
slightly  with  acetic  acid.    Add  an  equal  volume  of  alcohol  to  the  solution  or  as 
much  less  as  is  sufficient  to  cause  incipient  precipitation.     Stand  aside  for  15 
to  20  minutes.     Filter  off  the  precipitate,  wash  with  a  little  ether*   and  dry 
thoroughly.     Remove   from   filter   and   treat   in   a   test-tube   with   AgN03   and 
sufficient  NH4OH  to  dissolve  the  precipitate.     Place  the  test-tube  in  a  beaker 
of  boiling  water  and  thus  evaporate  to  dryness. 

The  formation  of  a  silver  mirror  indicates  the  presence  of  H2(C4H4O6). 

(3)  Acidify  10  ccm.  of  the  prepared  solution  with  HC1  and  boil  thoroughly. 
Into  the  hot  solution  pass  H2S  as  long  as  any  action  takes  place,  add  BaCl2  and 
filter.     Make  the  filtrate  distinctly  alkaline  with  Na2C03,  boil  thoroughly  and 
filter.    Make  filtrate  slightly  acid  with  acetic  acid  and  boil,  then  add  CaCl2  and 
stand   aside   for   5-10   minutes,   then   filter.     Make   filtrate   just   alkaline   with 
NH4OH.     White  precipitate  may  indicate  tartaric  acid.     Confirm  by  filtering 
and  drying  the  precipitate,  treating  with  AgN08  and  NH4OH  and  evaporating 
to  dryness  in  boiling  water  as  in  Test  (2). 

The  formation  of  a  silver  mirror  indicates  H2(C4H406). 

*  Ether  must  not  be  used  in  the  vicinity  of  any  flame. 


THE  ACID  ANALYSIS.  285 

DISCUSSION    OF    THE    TEST    FOR    OXALIC   ACID. 

If  the  solution  be  free  from  Ba,  Sr,  and  Pb  ions  (as  will  be  the  case  if 
the  solution  prepared  for  the  Acid  Analysis  be  used),  nothing  will  interfere  with 
this  test  for  oxalic  acid,  since  no  other  salts  of  calcium  which  are  more  in- 
soluble than  the  sulfate  will  be  precipitated  in  the  presence  of  acetic  acid. 
Since  oxalic  acid  is  easily  oxidized,  the  presence  of  oxidizing  agents  in  solution 
along  with  it  might  destroy  it  before  it  is  precipitated  as  the  calcium  salt. 


DISCUSSION   OF   THE    TESTS    FOR   TARTARIC    ACID. 

The  first  test  merely  indicates  the  presence  of  some  form  of  non-volatile 
organic  matter.  It  is  of  value  only  as  affording  evidence  as  to  whether  it  is 
necessary  to  make  the  more  elaborate  tests  since,  in  the  absence  of  carbonization, 
tartaric  acid  cannot  be  present. 

Test  (2)  effects  the  separation  of  tartaric  acid  from  other  acids  by  precipi- 
tating it  as  acid  potassium  tartrate  which  is  less  soluble  than  any  other  salt  of 
the  alkali  metals  with  the  exception  of  potassium  silicofluorid.  Metals  other 
than  the  alkalis  must  be  absent,  for  they  form  tartrates  more  insoluble  than  the 
acid  potassium  salt.  These  have  already  been  removed  in  making  the  prepared 
solution.  The  presence  of  K  ions  is  insured  by  the  addition  of  K2C03  and  an 
equal  volume  of  alcohol  is  added  to  the  solution  since  KH(C4H406)  is  much  less 
soluble  in  50  per  cent,  alcohol  than  in  water.  The  formation  of  a  precipitate 
does  not  necessarily  indicate  tartaric  acid  but,  if  this  acid  is  present,  it  will  be 
contained  in  the  precipitate  and  may  be  detected  by  the  reduction  of  AgN03. 
Since  alcohol  may  cause  this  same  reduction  it  is  necessary  to  remove  it  by 
washing  with  ether.  In  the  presence  of  a  slight  amount  of  NH4OH,  silver 
salts  are  readily  reduced  to  the  metal  by  organic  matter.  Tartaric  acid  accom- 
plishes this  reduction  on  warming  and  deposits  the  silver  as  a  brilliant  mirror 
on  the  sides  of  the  test-tube.  (For  further  discussion  see  pages  86  and  125.) 

Test  (3)  depends  upon  the  precipitation  of  tartaric  acid  as  the  calcium  salt 
and  the  subsequent  reduction  of  AgN03  as  confirmatory  evidence.  Inasmuch  as 
many  acids  (Groups  C,  D,  E;  see  page  276)  are  precipitated  by  the  calcium 
ion,  most  of  these  must  be  removed,  especially  those  which  will  react  with  AgN03 
or  reduce  it  to  the  metal.  Thorough  boiling  with  HC1  decomposes  and  removes 
H,S03  and  H2S203;  H2S  removes  the  arsenic  acids  completely  and  reduces 
dichromic  acid  to  the  basic  condition  and  BaCl2  precipitates  H,S04  and  H2SiFe 
By  boiling  the  filtrate  with  Na2C03,  chromium  and  the  excess  of  barium  are  pre- 
cipitated and  filtered  off.  The  filtrate  is  acidified  with  acetic  acid  and  boiled 
to  remove  H,S  and  C02.  The  addition  of  CaCl2  to  the  solution,  which  should 
be  but  faintly  acid  with  H(C2H302),  precipitates  oxalic  and  hydrofluoric  acids 
and  after  filtering  and  making  alkaline  with  NH4OH,  tartaric  acid  is  precipi- 
tated as  the  calcium  salt.  Certain  other  acids  may  be  present  in  the  precipitate, 
hence  the  indication  should  be  confirmed  by  the  formation  of  a  mirror  as  outlined 
in  Test  (2). 


286  QUALITATIVE    ANALYSIS. 

TESTS   FOR   PHOSPHORIC    ACID. 

(1)  Boil  some  of  the  prepared  solution,  or  dry  substance,  with  cone.  HN03 
as  long  as  brown  fumes  (denoting  oxidation)  are  given  off,  adding  more  acid, 
if  necessary,  to  complete  the  reaction.  Evaporate  to  dryness,  and  extract  the 
residue  by  boiling  with  dil.  HN03.  Filter,  if  necessary,  and  add  a  drop  or  two 
of  this  solution  to  5  ccm.  of  ammonium  molybdate,  (NH4)2Mo04. 

A  yellow  precipitate,  (NH4)3PO4.12MoO8,  indicates  H3PO4  (in  the  absence 
of  H3AsOj. 


(2)  If  arsenic  was  indicated  in  the  Basic  Analysis,  strongly  acidify  a  por- 
tion of  the  prepared  solution  with  HC1,  heat  to  boiling  and  pass  H2S  as  long 
as  a  precipitate  falls.  Filter,  boil  out  H2S  and  treat  as  in  Test  (1). 

A  yellow  precipitate,  (NH4)3PO4.12Mo03,  indicates  H3P04. 


THE  ACID  ANALYSIS.  287 

DISCUSSION   OF   THE   TESTS   FOR  PHOSPHORIC   ACID. 

The  formation  of  a  yellow  precipitate,  (NH4)sP04.12Mo03,  with  ammonium 
molybdate  (see  page  156),  is  a  reaction  characteristic  of  the  ortho-phosphoric 
acid,  but  not  of  any  other  acid  of  phosphorus.  The  presence  of  any  reducing 
agent  interferes  with  the  reaction  by  causing  a  reduction  of  the  molybdate, 
which  is  indicated  by  the  appearance  of  a  blue  color.  It  is  therefore  necessary 
to  evaporate  to  dryness  with  cone.  HN03  in  order  to  destroy  all  reducing  agents 
and  oxidize  any  lower  acids  of  phosphorus.  The  extraction  with  dilute  HN08 
causes  the  hydration  of  meta-  and  pyro-phosphoric  acid.  Thus,  by  this  prelimi- 
nary treatment  all  of  the  phosphorus  is  changed  to  ortho-phosphoric  acid. 

Arsenic  acid  with  ammonium  molybdate  gives  a  precipitate  closely  resem- 
bling that  with  phosphoric  acid  (see  page  164).  The  precipitate  due  to  H3P04 
will  come  down  in  the  cold  solution  after  a  few  minutes'  standing;  but  with 
H3As04  the  precipitate  cannot  be  brought  down  without  warming,  except  on 
long  standing.  This  difference  is  hardly  sufficient  to  differentiate  clearly;  hence 
arsenic  acid,  if  present,  must  be  removed  from  solution  before  preparing  for 
the  phosphoric  acid  test. 

Arsenic  present  in  a  form  other  than  H3As04  will  not  give  the  reaction  with 
molybdate;  hence  it  might  be  thought  that  this  would  not  interfere  with  the 
test.  But  under  the  treatment  prescribed  for  obtaining  phosphorus  in  the  form 
of  ortho-phosphoric  acid,  arsenic  originally  present  in  any  form  would  be 
oxidized  to  arsenic  acid.  Hence,  if  arsenic  is  indicated  in  the  Basic  Analysis,  it 
must  be  removed  as  in  Test  2  before  preparing  for  the  test  for  H3P04.  Arsenic 
present  in  any  form  is  precipitated  as  sulfid  by  this  treatment.  (For  further 
discussion  see  page  261.) 


288  QUALITATIVE    ANALYSIS. 

TEST    FOR    ARSENIC    ACID. 

Boil  a  portion  of  the  original  solution  with  Na2C08  to  precipitate  the  bases 
as  in  making1  the  prepared  solution,  and  filter.  Acidify  the  filtrate  to  break  up 
any  carbonate  and  neutralize  any  fixed  alkali,  boil  to  remove  C02,  make  alkaline 
with  NH4OH  and  add  NH4C1  and  MgCl2. 

The  formation  of  a  white  precipitate,  Mg(NH4)As04,  indicates  H3As04 
(in  the  absence  of  H3PO^  and  HJSiOJ. 

To  confirm,  filter  off  the  precipitate  and  wash  with  dilute  NH4OH.  Moisten 
the  precipitate  with  a  few  drops  of  AgN03. 

The  appearance  of  a  chocolate  color,  Ag3AsO4,  shown  at  once  or  on  the  addi- 
tion of  a  drop  of  acetic  acid,  indicates  H3As04. 


TEST   FOR   ARSENTOUS    ACID. 

From  the  solution  especially  prepared  as  directed  under  the  test  for  arsenic 
acid  (see  previous  test)  remove  H3As04  by  precipitation  as  Mg(NH4)As04  and 
filter.  Evaporate  the  filtrate  nearly  to  dryness  and  dissolve  in  3-5  ccm.  of  cone. 
HC1.  Add  3-5  drops  of  a  cone,  solution  of  SnCl2  in  cone.  HC1. 

The  appearance  of  a  dark  brown  precipitate,  metallic  As,  indicates  H3AsO3. 


THE  ACID  ANALYSIS.  289 

DISCUSSION   OF   THE   TEST   FOR  ARSENIC   ACID. 

The  test  for  the  acids  of  arsenic  must  be  made  on  a  portion  of  the  original 
solution  especially  prepared,  for  from  the  regularly  prepared  solution  all 
arsenic  has  been  precipitated  by  H2S. 

Since  all  the  salts  of  magnesium  except  the  arsenate  and  phosphate  (when 
formed  under  the  conditions  of  this  experiment)  are  soluble,  nothing  is  likely 
to  interfere  with  this  test.  The  ammonium  magnesium  arsenate  and  phosphate 
resemble  each  other  so  closely  as  to  be  indistinguishable  to  the  eye,  and  the 
differentiation  has  to  be  made  on  the  basis  of  the  transformation  into  the  silver 
salt,  which  with  arsenic  is  chocolate  brown,  with  phosphorus  pale  yellow. 

Silver  arsenate,  Ag3As04  is  readily  soluble  in  ammonium  hydroxid,  hence, 
after  washing  with  NH4OH,  it  may  be  necessary  to  neutralize  with  a  drop  of 
acetic  acid  to  bring  out  the  color. 

If  the  prepared  solution  is  yellow,  indicating  probably  the  presence  of 
chromates,  care  must  be  taken  to  wash  the  precipitated  Mg(NH4)As04  very 
thoroughly  with  NH4OH  until  every  trace  of  yellow  color  has  been  removed, 
for,  with  the  chromate  ion,  the  Ag  ion  gives  a  deep  red  precipitate  which  might 
be  confused  with  the  chocolate-colored  precipitate  due  to  arsenic  acid.  A  safer 
method  would  be  to  dissolve  the  precipitated  magnesium  salt  on  the  filter  paper 
with  very  dilute  HN03  and  wash  into  a  test-tube.  Then,  by  addition  of  dilute 
NH4OH,  it  may  be  reprecipitated  with  little  danger  of  inclusion. 

Dilute  NH4OH  is  used  to  wash  the  precipitate,  since  the  Mg(NH4)As04 
is  less  soluble  in  it  than  in  pure  water. 

It  might  appear  that  by  direct  addition  of  AgN03  to  a  neutral  solution 
H3As04  might  be  indicated  by  the  chocolate-colored  precipitate  at  once  without 
the  necessity  of  separating  it  as  the  ammonium  magnesium  salt.  Arsenic  acid 
is  far  less  frequently  met  with  in  analysis  than  most  other  inorganic  acids,  many 
of  which  give  precipitates  with  the  silver  ion  in  neutral  solutions  (if  not  in  the 
presence  of  HN03).  Since  Ag3As04  is  more  soluble  than  the  silver  salts  of 
many  other  acids,  H3As04  would  not  fall  until  after  the  other  acids  had  been 
removed  from  the  solution  by  precipitation;  thus  the  test  would  be  confused. 

If  soluble  silicates  are  present,  free  silicic  acid  may  be  precipitated  on  the 
addition  of  NH4C1  to  the  solution.  This  precipitate  is  due  to  hydrolysis  of  the 
ammonium  silicate  formed  by  metathesis  (see  page  135),  but  its  gelatinous  char- 
acter generally  serves  to  differentiate  it  from  the  crystalline  Mg(NH4)As04. 


DISCUSSION  OF   THE   TEST   FOR   ARSENIOUS   ACID. 

Since  the  solution  prepared  for  this  test  is  free  from  all  metals  except  arsenic 
and  antimony,  nothing  is  likely  to  interfere  with  this  test,  since  none  of  the  com- 
pounds of  antimony  will  give  a  similar  precipitate  with  SnCl2. 

If  arsenic  exists  in  the  prepared  solution,  and  be  not  present  in  the  form 
of  arsenic  acid  (which  will  be  removed  as  Mg(NH4)As04),  it  will  be  present  as 
arsenious  acid  or  its  equivalent. 

It  is  necessary  that  the  trivalent  arsenic  should  exist  in  the  solution  as 
arsenic  chlorid,  i.  e.,  acting  the  part  of  a  metallic  element,  and  not  as 
arsenious  acid.  Hence  it  is  absolutely  necessary  that  the  test  be  conducted  in 
the  presence  of  cone.  HC1.  (See  discussion  of  As03Hs,  page  159.) 

20 


290  QUALITATIVE    ANALYSIS. 

TEST   FOR   SULFUROUS   ACID. 

Acidify  a  portion  of  the  prepared  solution  with  acetic  acid  and  add  zinc 
acetate,  Zn(C2H302)2,  as  long  as  a  precipitate  falls.  Filter  off.  the  precipitate* 
and  test  the  nitrate  with  a  drop  of  Zn(C2H302)2  to  see  that  precipitation  is 
complete.  (Should  a  precipitate  fall,  add  more  zinc  acetate  solution  and  filter 
again.)  To  the  filtrate  add  BaCl2  in  considerable  excess  and  filter  until  a  clear 
solution  is  obtained.!  To  the  clear  filtrate  add  iodin  water  until  color  becomes 
permanent,  then  a  little  cone.  HC1  and  bleach  with  a  drop  or  two  of  SnCl2. 

White  precipitate,  BaS04,  indicates  H2S08. 

*  This  precipitate  may  be  tested  for  H2S  by  treating  with  H2SO4,  warming 
and  exposing  a  filter  paper  moistened  with  lead  acetate  solution  to  the  action  of  the 
fumes  evolved.  A  black  stain,  PbS,  indicates  the  presence  of  H2S. 

f  The  precipitate  obtained  here  may  be  dried  and  tested  for  H2SO4  by  heating 
before  the  blowpipe  with  Na2CO3  on  charcoal.  Eemove  the  melt,  place  on  a  silver 
coin  and  moisten  with  a  drop  of  water.  A  black  stain,  Ag2S,  indicates  the  presence 
of  H2SO4. 


TEST    FOR    THIOSULFURIC    ACID. 

Acidify  a  portion  of  the  prepared  solution  with  acetic  acid  and  add  zinc 
acetate,  Zn(C2H302)2,  as  long  as  a  precipitate  falls.  Filter  off  the  precipitate 
and  test  the  filtrate  with  a  drop  of  Zn(C2H302)2  to  see  that  precipitation  is 
complete.  (Should  a  precipitate  fall,  add  more  zinc  acetate  solution  and  filter 
again.)  To  the  filtrate  add  BaCl2  in  considerable  excess  and  filter  until  a  clear 
solution  is  obtained.  To  the  clear  filtrate  add  iodin  water  until  color  becomes 
permanent,  then  a  little  cone.  HC1  and  bleach  with  a  drop  or  two  of  SnCl24 
Filter  (if  necessary)  until  the  filtrate  becomes  clear,  add  bromin  water  until  the 
solution  becomes  permanently  colored  and  bleach  with  a  drop  or  two  of 
SnCl2. 

White  precipitate,  BaS04,  indicates  presence  of  H2S.,O3  (in  the  absence  of 
HSCN). 

%  The  directions  for  making  the  test  for  thiosulfuric  acid  up  to  this  point  are 
identical  with  the  test  for  sulfurous  acid  (see  above).  Time  may  be  saved,  there- 
fore, by  testing  for  thiosulfuric  acid  immediately  after  testing  for  sulfurous  acid 
by  adding  bromin  water  to  the  solution  (filtered  clear,  if  necessary)  that  was  used 
in  testing  for  sulfurous  acid  and  proceeding  as  directed. 


THE  ACID  ANALYSIS.  291 

DISCUSSION    OF    THE    TEST    FOR    SULFUROUS    ACID. 

The  test  for  sulfurous  acid  depends  upon  its  oxidation  by  iodin  to  sulfuric 
acid  (see  page  185)  and  its  subsequent  precipitation  as  the  barium  salt.  To 
prevent  a  similar  oxidation,  H2S  (if  present)  is  removed  as  ZnS  and  any 
H2S04  already  present  is  precipitated  as  BaS04.  If  on  the  addition  of  iodin 
water  to  the  solution  thus  prepared  a  second  precipitation  of  BaS04  falls,  it 
must  be  due  to  H2S03,  for  this  compound  only  is  oxidized  to  H2S04  by  iodin. 

Nothing  will  interfere  with  this  test  except  the  presence  of  oxidizing  agents 
which  will  act  on  sulfurous  acid  as  soon  as  the  solution  is  acidified.  Of  these 
dichromic  acid  will  be  precipitated  as  BaCr04  along  with  H2S04  and  thus 
removed  in  the  preliminary  treatment.  Of  the  other  common  oxidizing  acids 
(see  page  276)  the  barium  salts  are  all  soluble  and  may  be  separated  from  the 
sulfite  which  is  insoluble.  To  the  neutral  or  slightly  alkaline  solution  BaCl2  is 
added  as  long  as  precipitation  follows.  This  precipitate,  which  will  contain 
among  other  salts  BaS03,  is  filtered  off  and  treated  with  HC1  when  the  sulfite 
dissolves  and  may  be  tested  for  in  the  solution  as  directed. 

Bleaching  of  the  solution  with  SnCl2  (see  page  138)  renders  a  slight  precipi- 
tate more  plainly  visible.  It  may  be  omitted  if  the  precipitate  is  evident. 

By  testing  the  precipitate  brought  down  by  zinc  acetate  for  H2S  and  the 
first  precipitate  by  BaCl2  (before  the  addition  of  iodin  water)  for  H2S04  (as 
directed  in  the  footnotes  on  the  opposite  page),  these  acids  may  be  detected  also. 
If  thiosulfuric  acid  is  tested  for  in  the  same  solution  (see  next  test),  all  of  the 
four  sulfur  acids,  H2S,  H2S04,  H2S03  and  H2S208,  may  be  detected  in  the 
presence  of  each  other. 


DISCUSSION   OF  THE  TEST  FOR  THIOSULFURIC   ACID. 

This  test  depends  upon  the  fact  that,  when  treated  with  iodin,  thiosulfuric 
acid  is  oxidized  only  to  tetrathionic  acid,  the  barium  salt  of  which  is  soluble. 

2H2S203  -f  21  -»  H2S406  +  2HI 

The  stronger  action  of  bromin,  however,  oxidizes  tetrathionic  acid  to  sulfuric 
acid  (see  page  189)  and  in  the  presence  of  the  barium  ion  a  precipitate  of  BaS04 
is  obtained  which  in  this  case  indicates  H2S203. 

After  the  removal  of  other  sulfur  acids  (see  discussion  of  the  previous  test) 
nothing  will  interfere  with  the  detection  of  thiosulfuric  acid  except  the  presence 
of  oxidizing  agents  and  of  thiocyanic  acid.  The  former  may  be  separated 
exactly  as  indicated  in  the  discussion  of  the  test  for  sulfurous  acid;  for  barium 
thiosulfate,  like  barium  sulfite,  is  quite  insoluble  in  water. 

Thiocyanic  acid,  if  present,  will  be  oxidized  to  H2S04  by  bromin  and,  hence, 
must  be  removed.  This  is  most  satisfactorily  accomplished  before  beginning 
the  test  for  sulfurous  acid  by  faintly  acidifying  the  solution  with  HN03  and  add- 
ing AgN08  as  long  as  a  precipitate,  AgSCN  and  possibly  other  salts,  falls. 
The  nitrate  is  made  slightly  alkaline  with  NaOH  and  again  filtered,  after  which 
the  regular  treatment  for  H2S03  and  H2S203  is  to  be  followed. 


292  QUALITATIVE    ANALYSIS. 

TEST    FOR    FERROCYANIC    ACID. 

Acidify  a  portion  of  the  prepared  solution  with  HC1  and  add  a  drop  of 
FeCl3  solution. 

A  deep  blue  color,  Fe4(Fe  =  C6N6)3,  indicates  H4(Fe  =  C6NR). 


TEST  FOR  FERRICYANIC  ACID. 

Acidify  a  portion  of  the  prepared  solution  with  HC1  and  add  to  it  a  few 
drops  of  a  freshly  prepared  solution  of  FeS04. 

A  dark  blue  precipitate,  Fe3(Fe  =C6N6)2,  indicates  H3(Fe  =  C6N6). 


TESTS   FOR  THIOCYANIC  ACID. 

(1)  Acidify  a  portion  of  the  prepared  solution  with  HC1  and  add  FeCl3. 

A  deep  red  color,  Fe(SCN)3,  bleached  by  the  addition  of  HgCl2,  indicates 
HSCN. 

(2)  Acidify  a  portion  of  the  prepared  solution  with  HC1;   add  SnCl2,  if 
necessary,  to  hold  back  iodin  or  bromin  which  may  be  liberated  and  by  the  color 
produced  interfere  with  the  test  to  follow.     Boil,  collecting  the  volatile  thio- 
cyanic  acid  in  a  trap  containing  FeClj. 

The  formation  of  a  red  color  due  to  Fe(SCN)8  indicates  HSCN. 


THE  ACID  ANALYSIS.  293 

DISCUSSION   OF   THE   TEST  FOR   FERROCYANIC   ACID. 

Because  of  the  intensity  of  the  color  produced,  nothing  is  likely  to  interfere 
with  this  test.  When  only  a  small  amount  of  ferrocyanic  acid  is  present  in  a 
solution,  the  deep  blue  precipitate  may  be  so  sparsely  diffused  as  to  produce 
a  green  rather  than  a  blue  color,  especially  if  the  solution  be  colored  yellow 
by  an  excess  of  ferric  chlorid  present.  Under  these  circumstances,  if  the  indi- 
cation is  doubtful,  it  is  best  to  pour  the  solution  into  a  test-tube  and  allow  the 
precipitate  to  settle,  when  its  deep  blue  characteristic  appearance  will  suffi- 
ciently indicate  its  nature. 


DISCUSSION   OF   THE   TEST   FOR   FERRICYANIC   ACID. 

Nothing  is  likely  to  interfere  with  the  successful  working  of  this  reaction 
unless  the  presence  of  oxidizing  agents  should  carry  the  iron  from  the  ferrous 
to  the  ferric  condition.  Under  these  circumstances  it  is  better  to  reduce  oxidiz- 
ing agents  with  sodium  sulfite  in  acid  solution.  In  an  acid  solution  a  ferricyanid 
is  not  appreciably  reduced  to  a  ferroeyanid. 

Ferrous  salts  oxidize  so  quickly  in  the  air  that  it  is  essential  that  a  fresh 
solution  be  made  for  each  test  by  shaking  a  crystal  of  FeS04  with  a  few  cubic 
centimeters  of  water.  An  intense  dark  blue  color,  appearing  immediately,  alone 
indicates  the  presence  of  ferricyanic  acid;  for,  if  ferrocyanic  acid  is  present, 
under  the  conditions  of  the  experiment  a  light  blue  precipitate  will  form  which 
will  grow  visibly  darker  in  a  few  moments,  the  deepening  of  the  color,  and  the 
oxidation  of  the  ferrous  salts  causing  it,  being  hastened  by  vigorous  shaking  or 
pouring  the  solution  through  the  air. 


DISCUSSION   OF   THE    TESTS   FOR  THIOCYANIC   ACID. 

In  a  neutral  solution  acetic  acid  would  interfere  with  Test  (1)  by  the  pro- 
duction of  a  red  color  due  to  ferric  acetate.  Under  the  same  conditions  tartaric 
acid  and  similar  hydroxy  acids  tend  to  form  complex  compounds  which  remove 
ferric  ions  from  solution.  In  an  acid  solution  none  of  these  substances  interfere. 

lodids,  if  present,  may  interfere  with  the  test  for  traces  of  thiocyanic  acid 
in  the  way  already  indicated  (see  page  283).  If  oxidizing  agents  by  present 
in  solution,  iodin  may  be  set  free  in  quantity  when  the  solution  is  made  acid, 
and  therefore  may  interfere  decidedly.  This  interfering  action  may  be  avoided 
by  carefully  bleaching  out  the  solution  with  stannous  chlorid,  avoiding  an 
excess,  since  the  stannous  ion  will  reduce  ferric  salts  to  the  ferrous  condition 
and  thus  require  a  large  amount  of  ferric  chlorid  to  bring  out  the  red  color  due 
to  thiocyanic  acid. 

If  ferrocyanic  or  ferricyanic  acids  be  present  in  the  solution,  the  deep  blue 
color  due  to  these  reagents  may  completely  mask  the  red  color  due  to  thiocyanic 
acid.  Under  these  circumstances  it  is  best  to  distil  out  thiocyanic  acid  as  directed 
in  Test  (2). 


294  QUALITATIVE  ANALYSIS. 

TESTS    FOR    HYDROCYANIC   ACID. 

(1)  Make  a  portion  of  the  prepared  solution  alkaline  with  NaOH,  add  a  few 
drops  of  FeS04  solution  and  boil.     Add  a  single  drop  of  FeCl3  and  acidify 
with  HC1. 

A  deep  blue  color,  Fe4(Fe  =  C6N6)3,  indicates  HCN. 

(2)  Add  to  a  portion  of  the  prepared  solution  a  drop  or  two  of  yellow 
ammonium  sulfid  and  boil  for  a  few  minutes.     Cool,  acidify  with  HC1  and  add 
FeCl3. 

A  red  color,  Fe(SCN)3,  indicates  HCN. 

(3)  Make  a  portion  of  the  prepared  solution  alkaline  with  NaOH  and  add 
a  drop  of  a  dilute  solution  of  mercurous  nitrate,  (HgN03)2.     A  black  precipi- 
tate of  mercurous  oxid,  Hg20,  will  form. 

If,  on  shaking,  the  black  precipitate  disappears  and  a  gray  one,   due  to 
metallic  Hg,  appears  in  its  place,  it  is  an  indication  of  the  presence  of  HCN. 


TESTS  FOR  HYDROSULFURIC  ACID. 

(1)  Treat  a  portion  of  the  dry  substance,  or  the  prepared  solution,  with  HC1 
and  heat,  allowing  the  fumes  to  come  in  contact  with  a  filter  paper  moistened 
with  Pb(C2H302)2. 

A  black  stain,  PbS,  indicates  H2S. 

(2)  If  no  black  stain  appears  in   (1)   and  the  original  substance  remains 
undissolved,  fuse  it  with  NaOH  on  a  crucible  cover.    Place  the  melt  on  a  silver 
coin  and  moisten  with  a  few  drops  of  water. 

A  black  stain,  Ag2S,  indicates  H2S. 


THE  ACID  ANALYSIS.  295 

DISCUSSION    OF    THE    TESTS    FOR    HYDROCYANIC  ACID. 

Since  Test  (1)  depends  upon  the  transformation  of  the  simple  cyanid  ion,  CN, 
into  the  complex  ferrocyanid  ion,  (Fe  =  C6N6),  if  ferrocyanids  be  present  in  the 
original  solution,  the  result  of  the  test  is  obviously  worthless.  Ferricyanic  acid 
gives  a  deep  blue  color  with  ferrous  salts,  hence,  if  present,  it  will  interfere 
with  the  test.  Furthermore,  if  thiocyanic  acid  in  considerable  amount  and  traces 
only  of  hydrocyanic  acid  are  present,  the  intense  red  color  due  to  the  former  will 
probably  mask  the  slight  blue  tint  due  to  the  latter,  so  that  tests  cannot  be  made 
satisfactorily  on  the  solution.  For  similar  reasons,  the  second  method  is  no 
more  satisfactory  than  the  first  under  these  circumstances. 

Advantage  is  taken,  therefore,  of  the  weak  acidic  properties  of  HCN  to 
separate  it  from  the  other  stronger  cyanogen  acids  which  may  be  present.  The 
solution  is  made  exactly  neutral  with  HC1  or  NaOH,  and  dry  acid  sodium  car- 
bonate, NaHC03,  is  added  to  it.  Since  carbonic  acid  is  stronger  than  hydro- 
cyanic acid,  but  weaker  than  ferro-,  ferri,  or  thiocyanic  acids,  HCN  alone  will 
be  set  free  under  these  conditions.  On  gently  warming  the  solution  HCN  may 
be  distilled  out  and  collected  in  a  trap  containing  a  little  water.  More  satis- 
factory results  can  be  obtained  by  distilling  in  an  atmosphere  of  carbon  dioxid. 

Test  (2)  depends  upon  the  oxidation  of  HCN  to  HSCN  by  the  free  sulfur 
contained  in  the  solution  of  yellow  ammonium  sulfid.  (See  page  126.) 

Since  hydriodic  acid  is  oxidized  slowly  by  ferric  salts  with  the  liberation  of 
free  iodin,  a  slight  color  due  to  this  reaction  may  develop  on  standing  if  iodids 
are  present.  It  is  not  likely  to  interfere  seriously,  however,  as  the  action  of 
ferric  chlorid  would  take  place  in  slight  degree  only. 

Test  (3)  depends  on  the  reduction  of  mercurous  oxid  to  the  metal  through  the 
agency  of  the  cyanid,  as  the  following  equations  indicate: 

Hg20  +  2NaCN  -f  H20  -»  (HgCN)2  +  2NaOH 
(HgCN)2-»Hg(CN)2  +  Hg 


DISCUSSION    OF    THE    TESTS    FOR    HYDROSULFURIC    ACID. 

Certain  sulfids  are  not  readily  decomposed  by  HC1  and  therefore  do  not 
respond  satisfactorily  to  Test  (1).  When  fused  with  NaOH,  all  sulfids  are 
transposed  and  to  a  large  extent  oxidized  to  sulfates.  Sufficient  unoxidized 
Na2S  will  be  present,  however,  to  produce  a  stain  on  silver  due  to  the  formation 
of  Ag2S.  Oxygen  is  absorbed  during  the  reaction  according  to  the  equation : 

4Ag  +  2Na2S  +  02  +  2H20  ->  2Ag2S  +  4NaOH 


296  QUALITATIVE  ANALYSIS. 

TESTS   FOR   HYDRIODIC   ACID. 

(1)  Acidify  a  portion  of  the  prepared  solution  with  HN03,  add  one  to  two 
ccm.  of  CS2  and  treat  with  chlorin  water,  or  if  reducing  agents  are  present,  with 
K2Cr207,  shaking  the  mixture  vigorously  after  each  addition  of  the  reagent. 

The  appearance  of  a  purple  color  in  the  CS2  due  to  free  iodin  indicates  HI. 


(2)  Acidify  a  portion  of  the  prepared  solution  with  HN03,  add  KN02 
and  boil. 

The  evolution  of  purple  fumes  of  iodin  producing  blue  color  on  a  filter  paper 
moistened  with  starch  paste  indicates  HI. 


THE  ACID  ANALYSIS.  297 


DISCUSSION    OF    THE    TESTS    FOR    HYDRIODIC    ACID. 

Although  chlorin  water  will  liberate  both  bromin  and  iodin  from  their 
respective  acids  (see  page  200),  it  is  to  be  preferred  to  bromin  water,  which 
liberates  only  iodin  (see  page  208),  since  chlorin  does  not  color  CS2.  Since  all 
iodin  is  liberated  before  any  bromin  is  set  free,  with  careful  manipulation  no 
confusion  will  result  even  if  the  solution  contains  only  traces  of  iodid  in  the 
presence  of  much  bromid. 

Care  must  be  taken  to  avoid  the  addition  of  a  large  amount  of  chlorin  water 
at  the  start  since  an  excess  of  chlorin  will  unite  with  free  iodin  to  form  colorless 
iodin  chlorid  (see  page  211),  which  will  not  affect  the  CS2.  Hence  a  few  drops 
of  chlorin  water  should  be  added  at  first,  followed  by  larger  amounts  later 
if  an  indication  does  not  appear.  The  mixture  should  be  thoroughly  shaken 
after  each  addition. 

Certain  reducing  agents,  such  as  H2S,  H2S08,  H2S203,  HN02  especially, 
are  so  easily  oxidized  that  the  chlorin  is  entirely  absorbed  by  them  before  it 
can  set  free  iodin  from  hydriodic  acid.  If  they  be  present  in  the  solution  in 
relatively  large  proportion  there  may  be  no  indication  of  HI  even  after  the 
repeated  addition  of  the  necessarily  dilute  reagent,  chlorin  water.  If  these 
acids  are  suspected  or  known  to  be  present,  the  test  may  be  more  satisfactorily 
performed  by  the  use  of  K2Cr207  which  will  evidence  by  change  of  color  from 
yellow  to  green  that  it  is  being  reduced.  Sufficient  dichromate  is  added  to 
cause  the  appearance  of  a  yellow  tint  in  the  solution,  thus  indicating  an  excess 
of  the  reagent. 

Nitrous  acid  is  somewhat  soluble  in  CS2  forming  a  red-colored  solution.  If 
nitrites  be  present  when  Test  (1)  for  HI  is  made,  the  CS2  will  take  on  a  color 
which  might  be  confused  with  the  purple  tint  produced  by  iodin.  Since  iodin 
is  six  hundred  times  more  soluble  in  CS2  than  in  water,  the  color  due  to  iodin 
cannot  be  removed  by  shaking  with  water.  Nitrous  acid  is  about  as  readily 
soluble  in  one  as  the  other,  however,  and  the  color  due  to  it  can  be  removed 
by  pouring  off  the  solution  from  which  it  was  extracted,  washing  by  decanta- 
tion  and  shaking  vigorously  with  a  test-tube  of  water.  Under  these  conditions 
nitrous  acid  for  the  most  part  passes  into  the  water  and  the  color  leaves  the  CS2. 


298  QUALITATIVE  ANALYSIS. 

TESTS   FOR   HYDROBROMIC    ACID. 

(1)  To  a  portion  of  the  prepared  solution  acidified  with  HN03  add  1-2 
ccm.  of  CS2.  Treat  with  chlorin  water  and  shake  mixture  thoroughly. 

The  appearance  of  a  color  in  the  CS2,  red  or  yellow  according  to  the  con- 
centration of  the  free  bromin  present,  indicates  HBr. 

(2«)  Acidify  a  portion  of  the  prepared  solution  with  HN03  and  add  KMn04 
until  the  solution  is  colored  red  permanently.  Heat  to  boiling  and  collect  the 
volatilized  bromin  in  a  trap  containing  chloroform. 

A  reddish-brown  color  indicates  HBr. 


TESTS  FOR  HYDROCHLORIC  ACID. 

(1)  Acidify  a  portion  of  the  prepared  solution  with  HN03  and  add  AgN03. 

The  formation  of  a  curdy  white  precipitate,  AgCl  (readily  soluble  in 
NH4OH,  reprecipitated  on  the  addition  of  HN03,  darkening  in  the  sunlight), 
indicates  HCJ  (in  the  absence  of  HCN  and  HSCN). 

(2)*  Evaporate  some  of  the  prepared  solution  to  dryness,  add  solid  K2Cr20T 
and  5  ccm.  cone.  H2S04,  passing  the  fumes  through  a  trap  containing  a  drop  or 
two  of  water.  Rinse  out  the  trap  into  a  test-tube,  heat  to  boiling  if  necessary 
to  remove  free  bromin  and  iodin  and  add  a  grain  or  two  of  Na202. 

The  appearance  of  a  transient  blue  color,  HCr04,  at  once  or  on  acidifying 
with  H2S04,  denotes  the  presence  of  HC1. 


*  In  the  presence  of  chlorates  Test  (2)  is  prohibited  owing  to  the  explosive 
character  of  the  mixture  of  chlorates  and  cone.  H,SO4.  In  order  to  avoid  danger 
Test  (1)  for  HC1O3  (see  page  304)  should  be  made  before  making  this  test  for  HC1. 
If  chlorates  are  indicated,  the  separation  of  chlorids  may  be  effected  by  acidifying 
a  portion  of  the  prepared  solution  with  HNO3  and  adding  AgNO3  as  long  as  a  precipi- 
tate falls.  The  filtered  and  washed  precipitate  is  transferred  to  a  small  porcelain 
crucible,  covered  with  dilute  H2SO4  and  a  piece  of  Zn  laid  upon  it.  After  15-30 
minutes,  the  metathesis  represented  by  the  following  equation  will  have  taken  place 
and  the  decanted  solution  may  be  evaporated  to  dryness  and  tested  for  HC1  as  in  (2). 

2AgCl  +  Zn  — »  ZnCl2  +  2Ag 


THE  ACID  ANALYSIS.  299 

DISCUSSION    OF    THE    TESTS    FOR    HYDROBROMIC    ACID. 

Since  chlorin  water  liberates  both  bromin  and  iodin,  Test  (1)  cannot  be 
used  satisfactorily  when  both  HBr  and  HI  are  present,  since  iodin  is  liberated 
first  and  colors  the  CS2  purple.  Under  these  circumstances  permanganate  offers 
a  satisfactory  differentiation,  since  HBr  is  oxidized  to  water  and  free  bromin 
which  may  be  boiled  off  while  non-volatile  iodic  acid,  HI08  (see  page  211)  is 
formed  from  HI.  Care  should  be  taken  not  to  boil  the  solution  too  vigorously 
lest  the  heat  from  the  steam  vaporize  the  chloroform — boiling-point  61.5°. 


DISCUSSION  OF  THE  TESTS  FOR  HYDROCHLORIC  ACID. 

Under  the  conditions  of  the  experiment,  HC1,  HBr  and  HI  are  all  oxidized 
to  water  and  the  free  halogen.  Iodin  and  bromin  are  volatilized  but  chlorin 
unites  with  the  chromic  acid  anhydrid,  CrOs,  formed  by  the  dehydrating  ac- 
tion of  cone.  H2S04  on  K,Cr207,  to  produce  chloro-chromic  anhydrid,  Cr02Cl2 
(see  page  196).  Since  the  latter  compound  boils  at  116°  it  is  necessary 
to  heat  quite  strongly  to  drive  it  into  the  trap  where  it  reacts  with  water 
to  form  HC1  and  H2Cr207.  Inasmuch  as  H2Cr207  is  more  readily  detected 
than  HC1  under  the  conditions  of  the  experiment,  generally  the  former  is 
tested  for  by  Na202  and  its  presence  considered  as  indicating  the  presence  of  the 
latter  also.  Care  should  be  taken  to  prevent  the  liquid  in  the  test-tube  from 
being  carried  mechanically  into  the  trap  by  spattering,  as  this  would  give  a 
test  not  due  to  the  volatilized  Cr02Cl2. 

If  it 'is  necessary  to  make  tests  for  HC1,  HBr  and  HI  from  the  same  por- 
tion of  solution,  the  use  of  chlorin  or  bromin  water  has  to  be  avoided  since 
these  reagents  introduce  into  the  solution  two  of  the  acids  to  be  tested  for.  It 
is  necessary,  therefore,  to  detect  HI  according  to  the  second  test  as  given  under 
hydriodic  acid,  then  HBr  according  to  the  second  test  for  hydrobromic  acid 
and  finally  HC1  from  the  resulting  solution  evaporated  to  dryness. 


300  QUALITATIVE  ANALYSIS. 

TESTS    FOR    NITRIC    ACID. 

(1)  Saturate  a  portion  of  the  water  extract  of  the  original  substance  with 
FeS04  and  pour  cone.  H2S04  down  the  side  of  the  inclined  test-tube  in  such  a 
manner  that  the  acid  will  underlie  the  solution. 

The  formation  of  a  dark  ring,  due  to  (FeSOJ2.NO,  on  the  line  of  separation 
of  the  two  liquids  indicates  HN03  (in  absence  of  HN02). 


(2)  Acidify  a  portion  of  the  prepared  solution  with  H2S04  and  heat,  adding 
KI  drop  by  drop,  to  the  boiling  solution  until  all  iodin  and  oxids  of  nitrogen  are 
driven  off.  Then  make  the  solution  strongly  alkaline  with  NaOH  and  boil  out  all 
NH3  that  may  be  present.  Add  a  mixture  of  finely  divided  iron  and  zinc,  or  a 
little  Devarda's  alloy,  and  boil  again. 

The  evolution  of  NH3,  recognized  by  its  odor  or  action  on  red  litmus,  indi- 
cates the  presence  of  HNO3. 


(3)  Evaporate  a  portion  of  the  prepared  solution  almost  to  dryness  and 
cool.  Add  three  times  its  volume  of  cone.  H2S04  and  a  drop  of  the  brucin 
solution. 

A  red  color,  changing  to  yellow  and  possibly  green,  indicates  HN03  (in  the 
absence  of  HC103). 


THE  ACID  ANALYSIS.  301 

DISCUSSION   OF   THE   TESTS   FOR   NITRIC   ACID. 

Test  (1)  serves  for  the  detection  of  nitric  acid  in  the  presence  of  all  other 
acids  with  the  exceptions  of  nitrous  acid  (which  gives  the  same  test)  and  of  the 
interfering  substances  noted  below. 

Test  (2)  removes  HN02,  the  only  other  acid  that  will  give  the  same  reaction 
and,  therefore,  may  be  used  to  detect  nitric  acid  in  the  presence  of  any  other 
acids. 

Test  (3)  will  detect  nitric  acid  in  the  presence  of  nitrous  acid  and  of  any 
others  except  chloric  acid. 

lodids,  bromids,  ferrocyanids,  ferricyanids  and  oxidizing  agents  in  general 
interfere  with  Test  (1).  In  contact  with  cone.  H2S04,  hydrobromic  and  hydriodic 
acids  are  oxidized  with  the  liberation  of  iodin  and  bromin  which  produce  a 
similar  coloration.  The  ferrocyanid  and  ferricyanid  ions  form  blue  precipitates 
with  the  ferrous  ion,  making  the  test  very  indefinite.  Chromates,  dichromates 
and  permanganates  interfere  because  of  their  color. 

In  solutions  ordinarily  met  with  in  qualitative  analysis  ferrocyanids  and 
ferricyanids,  chromates,  and  permanganates  are  not  likely  to  be  present.  In 
many  kinds  of  work  bromids  and  iodids  occur  very  rarely.  Hence  the  test 
is  a  valuable  means  of  detecting  the  acids  of  nitrogen,  one  or  both,  and  is  the 
method  most  generally  used.  Although  the  interfering  acids  may  be  removed 
from  solution,  if  present  it  is  more  satisfactory  to  use  the  other  tests. 

Test  (1)  depends  upon  the  reduction  of  nitric  (or  nitrous)  acid  by  means 
of  ferrous  sulfate  and  the  interaction  of  the  reduction  product,  NO,  with  the 
same  reagent.  (See  page  151.) 

Test  (2)  depends  upon  the  fact  that  zinc  and  aluminum  are  acted  upon  by 
alkaline  hydroxids,  forming  sodium  zincate,  Na.,Zn02,  or  sodium  aluminate, 
Na3A103,  and  nascent  hydrogen  (see  pages  102  and  115).  The  iron  mixed  with 
the  zinc,  or  the  copper  in  the  alloy  forms  an  electrolytic  couple  within  the 
solution  and  aids  the  reaction.  Nascent  hydrogen  acting  on  nitric  acid  under 
these  conditions  reduces  it  to  ammonia.  The  presence  of  other  oxidizing  agents 
may  delay  but  cannot  prevent  the  formation  of  ammonia. 

Although  nitrous  acid  reacts  readily  with  hydriodic  acid  and  is  rapidly 
destroyed  (see  page  211)  as  in  Test  (2),  some  small  part  usually  decomposes  into 
nitric  acid  and  nitrogen  dioxid  according  to  the  regular  reaction  in  water 
solution  (see  page  148).  Hence  in  solutions  containing  nitrous  acid  only,  traces 
of  nitric  acid  are  frequently  indicated.  If  to  the  neutral  or  alkaline  solution  to 
be  tested  a  relatively  large  amount  of  cone.  H2S04  is  added,  this  formation  of 
nitric  acid  does  not  take  place.  Brucin  does  not  respond  to  nitrous  acid  and 
since  the  formation  of  nitric  from  nitrous  acid  is  prevented,  Test  (3)  becomes 
the  method  of  last  resort  for  the  detection  of  nitric  acid  in  the  presence  of 
nitrous  acid. 


302  QUALITATIVE    ANALYSIS. 

TESTS  FOR  NITROUS   ACID. 

(1)  Saturate  a  portion  of  the  water  extract  of  the  original  substance  with 
FeS04  and  make  slightly  acid  with  acetic  acid. 

A  dark  brown  color,  (FeS04)2.NO,  indicates  HN02. 

(2)  Acidify  a  portion  of  the  prepared  solution  and  add  a  few  drops  of 
phenylenediamin  in  HC1  solution. 

A  yellow  color  indicates  HNO2. 

(3)  Make  a  portion  of  the  prepared  solution  neutral  and  add  1  ccm.  of  the 
sulfanilic  acid  mixture. 

A  red  color  indicates  HNO2. 


TESTS    FOR    ACETIC    ACID. 

(1)   Treat  some  of  the  dry  substance  with  1-2  ccm.  of  the  ordinary  ethyl 
alcohol,  or  with  amyl  alcohol,  and  5  ccm.  of  cone.  H2S04  and  warm  (not  heat). 
The  evolution  of  the  characteristic  ethereal  odor  indicates  H(C2H302). 


(2)  To  the  neutral  solution  to  be  tested  add  FeCl3. 

The  appearance  of  a  deep  red  color  strongly  indicates  acetic  acid,  the  pres- 
ence of  which  may  be  confirmed  by  diluting  the  solution  abundantly  and  boiling, 
A  red  precipitate,  Fe(C2H3O2)(OH)2,  indicates  H(C2H302). 


THE  ACID  ANALYSIS.  303 

DISCUSSION   OF   TESTS  FOR  NITROUS  ACID. 

Test  (1)  for  nitrous  acid  is  very  similar  to  Test  (1)  for  nitric  acid  (see  page 
300),  the  difference  being  only  in  the  strength  of  the  acid  used  to  acidify  the 
solution.  Nitrous  acid  is  weaker  than  acetic  acid,  hence  is  liberated  by  it  and 
reacts  characteristically.  Nitric  acid,  which  reacts  similarly  to  nitrous  acid 
(see  pages  149  and  151),  is  much  stronger  than  acetic  acid,  hence  is  not  set  free. 
With  acetic  acid,  therefore,  the  dark  color  indicates  HNO,  alone;  but  when 
sulfuric  acid  is  used,  HN03  as  well  as  HN02  is  set  free  and  the  color  may  be 
due  to  either.  With  the  exception  of  HBr  (which  is  not  oxidized  by  HN02) 
the  same  substances  that  interfere  with  Test  (1)  for  nitric  acid  hinder  this  test 
for  nitrous  acid  also. 

The  colors  developed  in  Tests  (2)  and  (3)  are  due  to  characteristic  reactions 
of  nitrous  acid  (see  page  149).  Nothing  will  interfere  with  these  reactions. 
Should  an  indefinite  indication  be  obtained  because  of  very  dilute  solutions,  it 
may  be  advisable  to  make  a  blank  test  to  make  sure  of  the  indication.  Acidify 
about  20  ccm.  of  water  with  HC1,  add  a  few  drops  of  the  reagent  solution  and 
divide  equally  between  two  test-tubes.  To  one  add  the  solution  to  be  tested. 
The  development  of  a  tint  in  the  test-tube  to  which  the  unknown  solution  was 
added  indicates  the  presence  of  nitrous  acid  in  it. 


DISCUSSION    OF    TESTS    FOR    ACETIC    ACID. 

Test  (1)  depends  upon  the  formation  of  ethyl  or  amyl  acetate  by  the  union 
of  acetic  acid  with  the  corresponding  alcohol.  Alcohols  are  organic  bases  — 
hydroxids  —  and  in  the  reaction  with  acids  water  is  set  free  just  as  with  inor- 
ganic bases. 


25  H(C2H302)  ->  (C2H5)(C2H302)  +  H20 

ethyl  alcohol  ethyl  acetate 

(C5HU)OH  +  H(CJE80,)  ->  (C5H  )  (C2H302)  +  H20 

amyl  alcohol  amyl  acetate 

Concentrated  H2S04  is  added  to  liberate  acetic  acid  from  its  salts  and  to 
unite  with  the  water  formed  during  the  reaction  and  prevent  reversal  (see 
page  23). 

It  is  necessary  that  the  characteristic  odor  of  ethyl  and  amyl  acetates  be- 
come familiar  to  the  observer  in  order  that  he  may  differentiate  it  clearly  from 
the  odor  of  ethyl  or  amyl  alcohol.  If  necessary  a  blank  test  may  be  made. 

Test  (2)  depends  upon  the  hydrolysis  of  ferric  acetate  in  a  solution  slightly 
acid  with  acetic  acid.  In  order  to  keep  ferric  salts  in  solution  it  is  necessary 
to  add  a  little  free  acid.  When  a  ferric  salt  is  added  to  a  neutral  solution  of  an 
acetate,  therefore,  the  free  acicl  present,  usually  HC1,  liberates  a  little  acetic 
acid  and  forms  a  corresponding  amount  of  chlorid.  The  amount  of  free  acid 
is  therefore  the  determining  factor  in  the  success  of  this  test.  Care  must  be 
taken  that  the  solution  before  the  addition  of  the  ferric  salt  is  not  alkaline  but 
distinctly  neutral.  If  alkaline,  ferric  hydroxid  may  be  precipitated  and  fur- 
nish no  indication  at  all  of  the  presence  of  acetic  acid.  If  the  solution  be  acid 
ferric  acetate  may  not  be  formed  at  all  and  so  not  give  the  red  color  to  the 
solution.  At  any  rate,  if  the  solution  is  too  acid,  the  precipitation  of  basic 
ferric  acetate  would  be  prevented. 

Other  organic  acids,  not  included  in  this  manual  and  rarely  met  with  in 
analysis,  when  similarly  treated,  separate  basic  ferric  salts  indistinguishable 
from  the  basic  acetate. 


304  QUALITATIVE    ANALYSIS. 

TESTS    FOR    CYANIC    ACID. 

(1)  Neutralize   a  portion   of  the   prepared  solution   and   concentrate   to   a 
small  bulk.    Add  1  to  2  ccm.  of  a  strong  solution  of  cobalt  acetate,  Co(CfLt02)n. 

A  blue  coloration  due  to  K2(Co04C4NJ,  appearing  at  once  or  on  the  addi- 
tion of  an  equal  volume  of  alcohol,  indicates  HOCN  (in  the  absence  of  HSCN). 

(2)  If  not  already  so,  make  a  portion  of  the  prepared  solution  alkaline 
with  NaOH  and  boil  until  NH3  (if  present)  is  entirely  removed.     Cool,  acidify 
the  solution  with  H2S04  and  test  the  gas  evolved  for  C02,  liberated  on  acidifying 
a  solution  containing  a  cyanate.    Boil  until  all  H2S  (if  any  be  present)  has  been 
driven  off  along  with  CO,.     Make  the  solution  alkaline  with  NaOH  and  boil, 
testing  the  gas  evolved  with  red  litmus  paper  or  mercurous  nitrate  on  a  filter 
paper. 

The  turning  blue  of  the  litmus  paper  or  the  production  of  a  black  stain  with 
mercurous  nitrate  is  due  to  NH3  set  free  by  the  decomposition  of  HOCN. 


TEST  FOR  HYPOCHLOROUS  ACID. 

Make  a  portion  of  the  aqueous  extract  of  the  original  substance  strongly 
alkaline  with  Na2C03  and  add  3-5  drops  of  H(C2H302),  shaking  thoroughly 
after  each  drop.  Care  should  be  taken  not  to  add  acetic  acid  in  sufficient  quan- 
tity to  produce  a  marked  effervescence  or  to  make  the  solution  acid.  Insert  a 
piece  of  litmus  paper. 

Bleaching  of  the  litmus  indicates  HOC1  (in  the  absence  of  HMnOt,  which 
may  be  recognized  by  its  color). 

If  HOC1  is  present,  it  may  be  recognized  by  the  characteristic  odor  of  this 
solution — like  that  of  "  bleaching  salt." 


TESTS  FOR   CHLORIC  ACID. 

(1)  To  a  small  portion  of  the  dry  substance   (caution!— use  a  few  grains 
only)  add  cone.  H2S04  and  warm  gently. 

A  slight  explosion  or  crackling  due  to  C102  indicates  HC108. 

(2)  Acidify  a  small  portion  of  the  prepared  solution  with  HN03,  heat  to  boil- 
ing and  add  AgN03  until  precipitation  ceases.     Shake  thoroughly  to  coagulate 
the  precipitate,  filter  and  to  the  clear  filtrate  add  solid  Na2S03  until  the  odor  of 
S02  is  apparent. 

A  white  precipitate,  AgCl  (insoluble  in  boiling  water  to  which  a  little  HNOa 
is  added,  but  readily  soluble  in  NH4OH),  indicates  HC108. 


THE  ACID  ANALYSIS.  305 


DISCUSSION    OF    THE    TESTS    FOR    CYANIC    ACID. 

Since  the  test  with  cobalt  acetate  (see  page  128)  is  not  given  by  cyanates  in 
the  presence  of  cyanids,  these  are  removed  (if  present)  by  passing  C02  through 
the  neutral  solution  for  from  15-90  minutes,  the  time  necessary  being  dependent 
upon  the  amount  of  cyanid  present.  Under  these  conditions  hydrocyanic  acid 
is  set  free  (see  page  295)  and  swept  from  the  solution  by  the  carbon  dioxid, 
while  the  cyanate  is  not  appreciably  affected. 

Since  thiocyanates  give  the  test  with  cobalt  acetate  as  well  as  the  cyanates, 
in  the  presence  of  thiocyanates  the  indication  is  equivocal  and  resort  must  be 
had  to  the  second  test.  Any  ammonium  salts  present  in  solution  will  be  broken 
up  and  removed  on  boiling  the  alkaline  solution.  On  acidifying,  the  cyanate  is 
decomposed  into  NH3  which  unites  with  the  excess  of  acid  and  remains  in  solu- 
tion, while  C02  is  driven  off.  Of  course  the  presence  of  carbonates  would  lib- 
erate C02  so  that  its  detection  would  not  indicate  the  presence  of  cyanates  ex- 
cept in  the  absence  of  carbonates.  In  any  case  it  is  better  to  confirm  by  testing 
for  NH3  on  making  the  solution  alkaline.  Mercurous  nitrate  can  be  used  with 
advantage  to  detect  the  presence  of  NH3  (see  page  110),  since  all  H2S  (which 
also  produces  a  black  stain)  should  be  removed  by  boiling  the  acid  solution 
until  it  gives  no  stain  on  lead  acetate  paper. 


DISCUSSION  OF  THE   TEST  FOR  HYPOCHLOROUS  ACID. 

All  hypochlorites  are  soluble  and  will  therefore  be  found  in  the  aqueous  ex- 
tract. Silver  hypochlorite  decomposes,  however,  into  silver  chlorid  and  chlo- 
rate (see  page  205).  Consequently,  unless  a  precipitate  is  obtained  with  AgN03 
in  the  preliminary  testing  for  acids,  hypochlorites  are  not  present. 

The  test  depends  upon  the  fact  that,  HOC1  being  very  weak,  is  set  free 
from  its  salts  by  the  carbonic  acid  liberated  by  the  interaction  of  acetic  acid 
and  sodium  carbonate,  even  though  the  solution  is  alkaline  in  reaction  due  to  the 
presence  of  acid  carbonates  of  the  alkali  metals  (see  page  204). 


DISCUSSION    OF    THE    TESTS    FOR    CHLORIC    ACID. 

Test  (1)  depends  on  the  formation  of  C102  which  decomposes  explosively 
into  its  elements.  Chloric  acid  is  at  first  set  free  and  then  breaks  down,  accord- 
ing to  the  equation : 

3HC103  ->  HC104  +  H20  -f  2C102  -*  C12  +  202 

Care  must  be  taken  to  use  only  a  few  grains  of  the  substance  to  be  tested 
as  serious  explosions  result  from  the  interaction  of  chlorates  and  sulfuric  acid 
in  too  large  amounts. 

Test  (2)  depends  upon  the  reduction  of  HC103  to  HC1  by  sulfurous  acid. 
Chlorids  and  hypochlorites  are  removed  from  the  solution  by  the  addition  of 
the  Ag  ion  in  excess.  After  filtering,  the  sulfite  is  introduced.  If  a  further 
precipitate  of  AgCl  falls,  it  is  due  to  HC108,  for  HC104  is  not  decomposed  by 
ordinary  reducing  agents.  White  Ag,S04  or  Ag2S03  may  be  precipitated  from 
the  concentrated  solution,  but  either  will  dissolve  readily  when  boiled  with  water 
containing  a  little  HN03. 
21 


306  QUALITATIVE    ANALYSIS. 

TEST  FOR  PERCHLORIC  ACID. 

Acidify  a  small  portion  of  the  prepared  solution  with  HN03,  heat  to  boil- 
ing and  add  AgNOs  until  precipitation  ceases.  Shake  thoroughly  to  coagulate 
the  precipitate,  filter  and  to  the  clear  filtrate  add  solid  Na2S03  until  the  odor  of 
S02  is  apparent.*  Filter,  if  necessary,  and  boil  the  clear  filtrate  for  a  few 
moments.  Cool,  make  alkaline  with  NaOH,  evaporate  to  dryness  and  ignite 
strongly.  Cool,  extract  by  boiling  with  dilute  HN03  and  add  a  drop  or  two  of 
AgN03  if  a  precipitate  has  not  already  formed. 

A  white  precipitate  of  AgCl  (soluble  in  NH4OH  and  darkening  in  sunlight) 
indicates  HC1O4. 

*  The  directions  for  making  the  test  for  perchloric  acid  up  to  this  point  are 
identical  with  the  test  for  chloric  acid  (see  page  304).  Time  may  be  saved,  therefore, 
by  testing  for  perchloric  acid  immediately  after  testing  for  chloric  acid  by  con- 
tinuing the  treatment  as  further  directed. 


TEST   FOR   PERMANGANIC   ACID. 

If  the  aqueous  extract  of  the  original  substance  is  colored  red,  acidify  a 
portion  with  H2S04  and  add  Na2S03. 

Bleaching  of  the  red  color  indicates  the  presence  of  HMn04. 


TESTS   FOR   HYDROGEN   PEROXID. 

Treat  a  portion  of  the  finely  powdered  substance  with  dil.  H2S04,  warm 
gently  and  stand  aside  for  10-15  minutes,  shaking  vigorously  from  time  to 
time.  Filter  clear  and  use  the  filtrate  in  making  the  following  tests.  If  the 
original  substance  is  already  in  solution,  acidify  a  portion  with  H2S04,  filter  if 
necessary  and  use  this  filtrate  in  the  following  tests. 

(1)  To  one  portion  of  the  solution  prepared  as  above  directed  add  a  drop 
or  two  of  titanium  sulfate,  Ti(S04)2. 

The  appearance  of  a  color  varying  from  yellow  to  orange-brown  indi- 
cates H2O2. 

(2)  To   a  very  dilute   solution   made  by   adding   a   single   drop   of   FeCl3 
and  a  single  drop  of  freshly  prepared  K3(Fe  =  C6N6)   solution  to  10  ccm.  of 
water  acidified  with  H2S04,  add  a  few  drops  of  the  solution  prepared  as  above 
directed. 

The  appearance  of  a  green  or  blue  color  due  to  Fe4(Fe  =  C6N6)3  indi- 
cates H2O2. 


THE  ACID  ANALYSIS.  307 


DISCUSSION    OF    THE    TEST    FOR   PERCHLORIC    ACID. 

After  the  decomposition  of  hypochlorites  and  chlorates  into  chlorids  and 
the  removal  of  the  latter  as  AgCl,  the  solution  can  contain  no  other  acid  of 
chlorin  than  HC104.  By  evaporation  and  ignition  of  the  resulting  nitrate,  the 
perchlorate  is  decomposed  into  the  chlorid  and  oxygen.  By  the  treatment  with 
NaOH,  any  excess  of  AgN03  that  is  present  will  be  precipitated  as  the  hydroxid. 
This  decomposes  at  once  into  the  oxid  and  this  on  heating  into  metallic  silver. 
By  boiling  with  HN03,  AgN03  will  be  formed  again  and  a  precipitate  of  AgCl 
under  these  circumstances  will  indicate  HC104.  If  no  precipitate  is  seen,  it  is 
safest  to  add  a  few  drops  of  AgN03  to  be  certain  that  the  Ag  ion  is  present. 


DISCUSSION    OF    THE    TEST    FOR    PERMANGANIC  ACID. 

If  permanganic  acid  is  present  during  the  Basic  Analysis,  it.  will  be  reduced 
by  H2S  to  the  divalent  basic  condition  and  detected  in  Group  III.  If  man- 
ganese is  not  detected  in  the  Basic  Analysis,  it  is  unnecessary  to  test  for  the 
presence  of  HMn04  among  the  acids. 


DISCUSSION    OF    THE    TESTS    FOR    HYDROGEN    PEROXID. 

Of  the  commoner  peroxids  (salts  of  H202)  Na202  and  K202  are  soluble  in 
water.  Others  are  more  or  less  soluble  in  acids,  but  require  careful  treatment 
in  order  not  to  decompose  the  H202  formed.  For  this  reason  heating  is  to 
be  avoided. 

Test  (1)  depends  on  the  oxidation  of  titanium  from  the  tetravalent  to  the 
hexavalent  condition  which  is  characterized  by  the  deep  orange-red  color. 

Test  (2)  depends  upon  the  interaction  of  hydrogen  peroxid  and  ferricyanic 
acid  by  virtue  of  which  both  are  reduced,  the  former  to  water,  the  latter  to 
ferrocyanic  acid.  This  last  substance  reacts  with  the  ferric  salt  to  produce 
Prussian  blue.  Oxygen  is  liberated  by  the  interaction  of  H202  with 
H8(Fe  =  CaN3)  as  the  equation  shows: 

2H3(Fe  -  C6NG)  +  H202  ->  2H4(Fe  -  C6Nfl)  +  02 

Test  (2)  is  obviously  not  applicable  in  the  presence  of  ferrocyanic  acid  or  of 
large  amounts  of  thiocyanic  acid. 


308  QUALITATIVE    ANALYSIS. 

ANALYSIS  OF  ORGANIC  SUBSTANCES. 

(1)  Carbon  and  Hydrogen.     Mix  a  small  amount  of  the  substance  to  be 
tested   with   about   five   times  its  weight   of   finely   powdered   CuO    previously 
ignited,  and  heat  in  a  test-tube,  passing  the  fumes  evolved  through  a  trap  con- 
taining "  lime  water." 

A  white  precipitate,  CaC03,  indicates  the  presence  of  C.  Moisture,  H20, 
deposited  on  the  cooler  parts  of  the  test-tube  indicates  H. 

(2)  Nitrogen.     Ignite  a  portion  of  the  substance  with  a  small  piece  of 
metallic  sodium  in  a  tube  of  hard  glass.     Drop  the  tube  while  still  hot  into  a 
beaker  containing  20  or  30  ccm.  of  water.     Boil  and  filter.     Add  to  the  filtrate 
a  small  crystal  of  FeS04,  and  a  drop  of  FeCl3.     Boil  for  a  few  minutes  and 
acidify. 

A  deep  blue  color,  Fe4(Fe  =  C8N6)3,  indicates  N. 

(3)  Sulfur.     Ignite  a  portion  of  the  substance  with  metallic  sodium  as  in 
(2).    Boil  out  with  water,  filter  and  divide  the  solution. 

(a)  Into  one  portion  drop  a  bright  silver  coin  and  allow  to  stand  for  some 
time. 

Light  brown  stain  of  Ag2S  indicates  the  presence  of  S. 

(5)  To  the  second  portion  add  a  few  drops  of  a  solution  of  sodium  nitro- 
ferricyanid,  Na2Fe(CN)5NO. 

A  purple  color  indicates  S. 

(c)   Slightly  acidify  a  third  portion  with  H(C2H302)  and  add  Pb(C2H302)2. 

Black  precipitate,  PbS,  indicates  S. 

(4)  Chlorin,  Bromin  and  lodin.     Ignite  a  portion  of  the  substance  in  a 
test-tube  with  lime.     Cool,  boil  out  with  water  and  filter. 

(a)  To  a  small  portion  of  the  filtrate  add  HN03  and  AgN08. 
A  precipitate,  AgCl,  AgBr  or  Agl,  indicates  01,  Br  or  I. 

(b)  In  case  a  precipitate  is  obtained  in  (a),  use  the  remainder  of  the  solu- 
tion to  determine  what  members  of  the  halogen  group  are  present,  according 
to  the  methods  outlined  on  page  299. 

(5)  Phosphorus.    Boil  a  portion  of  the  substance  with  cone.  HN03  as  long 
as   the   evolution   of   brown   fumes   indicates   that    oxidation   is   taking   place. 
Evaporate  to  dryness,  cool,  and  extract  with  a  little  dil.  HN03.    Filter  and  add 
a  drop  or  two  of  this  solution  to  five  ccm.  of  (NH4)2Mo04. 

Yellow  precipitate,  (NH4)3P04.12Mo03,  indicates  P. 


THE  ACID  ANALYSIS.  309 


DISCUSSION  OF  THE  ANALYSIS  OF  ORGANIC  SUBSTANCES. 

The  tests  given  in  the  regular  analysis  are  tests  for  ions,  or  in  a  few  cases, 
for  free  elements  (Cl,  Br,  I,  etc.).  Since  organic  compounds  do  not  generally 
ionize,  they  will  not  respond  to  these  tests.  In  order  to  determine  the  constitu- 
ents of  these  bodies,  it  is  necessary  to  decompose  them  into  substances  which 
will  readily  ionize  or  substances  for  which  definite  tests  are  known.  Thus,  C, 
H,  and  P  are  recognized  by  their  oxidation  products,  C02,  H20,  and  H8P04 
Sulfur  contained  in  organic  compounds  is  changed  to  sodium  sulfid  which  gives 
tests  for  the  sulfid  ion.  Similarly  the  halogens  heated  with  lime  are  changed 
into  compounds  which  readily  ionize. 

The  test  for  carbon  and  hydrogen  depends  upon  the  oxidation  of  the  sub- 
stance with  copper  oxid  and  the  formation  of  C02  and  water.  Since  all 
organic  compounds  contain  carbon  and  practically  all  of  them  contain  hydrogen, 
this  test  is  of  significance  only  as  determining  whether  or  not  a  substance  is  of 
organic  nature. 

The  test  for  nitrogen  depends  upon  the  fact  that  whenever  organic  material 
containing  both  carbon  and  nitrogen  is  heated  with  an  alkali  metal,  a  cyanid  is 
formed.  When  an  alkali  cyanid  is  boiled  with  a  ferrous  salt,  potassium  ferro- 
cyanid  is  formed  (see  page  127),  which  with  the  ferric  ion  gives  Prussian  blue. 

Free  sulfur  or  sulfur  contained  in  organic  compounds,  when  heated  with 
metallic  sodium,  forms  sodium  sulfid.  According  to  the  amount  of  sulfur  pres- 
ent in  the  original  substance,  indications  may  be  obtained  with  the  various 
reagents  mentioned.  The  test  with  the  silver  coin  is  more  delicate  than  the 
other  two  and  will  indicate  a  trace  of  sulfur  when  no  evidence  would  be  obtained 
by  the  second  or  third  tests.  Sodium  nitro-ferricyanid  will  respond  to  the 
presence  of  appreciable  amounts  of  sulphur,  while  lead  sulfid  will  not  be  precipi- 
tated except  in  the  presence  of  considerable  amounts  of  sodium  sulfid  due  to  a 
high  percentage  of  sulfur  in  the  original  substance. 

In  the  decomposition  of  the  organic  substance  due  to  heating,  any  halogens 
present  will  unite  with  the  lime  to  form  calcium  compounds,  which  will  then 
give  the  ordinary  tests  for  the  halogen  ions. 

By  heating  with  cone,  nitric  acid  an  organic  substance  may  be  completely 
oxidized,  any  phosphorus  present  being  changed  to  phosphoric  acid,  which  will 
give  any  of  the  regular  tests. 


APPENDIX 


REAGENTS    IN    SOLUTION. 


Acids. 


Hydrochloric,  HC1,  concentrated— 

Sp.  Gr.  1.190  37.2%  HC1 

1.195  38.2 

1.200  39.1 

dilute 


Nitric,  HN03,  concentrated— 

Sp.  Gr.  1.42  69.8%  HN03 

dilute 
Sulfuric,  H2S04,  concentrated— 

Sp.  Gr.  1.839  95    %  H2S04 

1.840  95.6 

1.841  97 
dilute 


Acetic,  H(C2H302),  dilute 
Hydrofluoric,  HF,  concentrated,  40% 


12.1N 
12.5N 
12.9N 

4  N 
4  N 
4  N 

15.7N 


35.7N 

36    N 

36.5N 

4   N 

4  N 
4  N 
4  N 


Volume  to  be  diluted  to  1  litre. 


330  ccm.  of  37.2%  acid. 
320  ccm.  of  38.2%  acid. 
311  ccm.  of  39.1%  acid. 


254  ccm.  of  69.8%  acid. 


112  ccm.  of  95  %  acid. 
Ill  ccm.  of  95.6%  acid. 
110  ccm.  of  97  %  acid. 
277  ccm.  of  80  %  acid. 


Alkalies. 

Ammonium  hydroxid,  NH4OH,  concentrated — 
Sp.  Gr.  0.90  28.3%  NH3 

dilute 


15   N 
4   N    266  ccni.  of  28.3%  sol. 


Grams  of 

Substance 

in  One 

Litre  of 

Solution.  § 

Potassium  hydroxid,  KOH.  4  N  281 

The  material  used  for  analytical 
purposes  contains  about20%  water. 
Hence  the  amount  necessary  for  a 
solution  four  times  normal  is  (56.2  X 
4)  X  5/4  =  281. 

§  Not  in  one  litre  of  water.  The  substance  should  be  dissolved  in  a  quantity  of 
distilled  water  sufficient  for  the  purpose  and  the  resulting  solution  then  diluted  to 
1000  ccm.  with  distilled  water. 

313 


314  APPENDIX. 

Sodium  hydroxid,  NaOH.  4  N 

The  material  used  for  analytical 
purposes  contains  about  10%  water. 
Hence  the  amount  necessary  for  a 
solution  four  times  normal  is  (40  X 
4)  X  10/9  =  177.7. 

Other  Reagents  in  Solution  or  in.  Liquid  Form. 

Alcohol — aniyl 

— ethyl,  or  ordinary 

— methyl,  or  wood 
Aluminum  chlorid,  A1C13.6H20  N/2 

—alum,  KA1(SOJ2.12H20  N/2 

Ammonium  acetate,  NH4(C2H302)  4   N 

— carbonate,  (NHJ2C03,  (see  page  318) 

about  4   N 

—chlorid,  NH4C1  4    N 

— molybdate,  (NH4)Mo04,  (see  page  318)  N 

— oxalate,   (NH4)2C204.H20,  about  N/2 

— sulfid    (colorless),    (NHJ2S,    (see 

page  318),  about  4   N 

— polysulfid   (yellow),   (NHJ2S  X  S, 

(see  page  318) 
Aqua  regia,  (see  page  43) 
Barium  chlorid,  BaCl2.2H20  N 

—hydroxid,  Ba(OH)2.8H20,  about  N/3 

Bromin  water,  about  N/2 

Brucin,  (see  page  319) 
Calcium  chlorid,  CaCl2  N  55.6 

—hydroxid,  Ca(OH)2,  "lime  water," 

about  N/22  sat.  so!4 

— sulfate,  CaS04.2H20,  about  N/33  sat.  so!4 

Carbon  disulfid,  CS2 

Chlorin  water,  about  •  N/5  sat.  sol. 

Chloroform 
Cobalt  acetate,  Co  ( C2H302)  .4H20  N  124.5 

—nitrate,  Co(N03)2.6H20  N  147.8 

Ether 

Ferric  chlorid,  FeCl3*  ^     N  54.1 

lodin  water,  (see  page  319),  about  N/4 

Lead  acetate,  Pb(C2H302).3H20*  N  189.5 

"  Magnesia  mixture,"  (see  page  319)  N 

*  Should  contain  a  little  free  acid. 

$  Saturated  solutions  should  be  made  by  allowing  the  solvent  to  stand  in  contact 
with  the  solute  for  considerable  time,  the  mixture  being  shaken  occasionally.  After 
settling  thoroughly,  the  clear  solution  should  be  decanted  or  siphoned  off  into  another 
bottle  for  use. 


EEAGENTS   IN   SOLUTION.  315 

Manganese     chlorid,     MnCl2.4H20,     (see 
page  319) 

Mercuric  chlorid,   HgCl2,   about  N/2  sat.  so!4 

Mercurous  nitrate,   (HgN03)2,   (see  page 

319)  N/2  135.7 

Phenylenediamin,  (see  page  319) 

Potassium  chromate,  K2Cr04  N  97.3 

— cyanid,  KCN  N  65.2 

— dichromate,  K2O207  N  73.8 

— ferricyanid,  K3(Fe  =  C6N6)  N  109.8 

— ferrocyanid,  K4(Fe-C6Ne)  N  105.7 

— iodid,  KI  N  166.2 

—nitrite,  KN02  N  85.2 

— permanganate,  KMn04  N  79.1 

— thiocyanate,  KSCN  N  97.2 

Silver  nitrate,  AgN03  N  169.9 

Sodium  acetate,  Na(C2H302)  4   N  328 

— carbonate  (dry),  Na2C03  or  about      4    N  sat.  sol.;}: 

—carbonate  (crystals),  Na3C03.10H20     4   N  572 

— nitroprussid,       Na2(Fe  =  C5N5NO) 

.2H20,  5  gr.  in  100  ccm.  H20 

—phosphate  (acid),  Na2HP04.12H20          N  119 

— stannite,   Na2Sn02,    (see   foot-note, 

page  254) 

— sulfid,  Na2S,  10  gr.  in  100  ccm.  H20 
— sulfite,  Na2S03.8H20  N  135 

Sodium  thiosulfate,  Na2S203.5H20  N  124 

Stannic  chlorid,  SnCl4,   (see  page  319)  N/2  32.7 

Stannous     chlorid,     SnCl2,     concentrated 
used  in  Bettendorff  test  for  arsen- 
ious  acid,  (see  page  319) 
—chlorid,  SnCl2.2H20,  (see  page  319)         N/2  56.5 

Starch  solution,  (see  page  319) 

— iodid  solution,  (see  page  320) 

Sulfanilic  acid  mixture,  (see  page  320) 

Titanium  sulfate,  Ti(S04)2,  10  grams  to 
100  ccm.  H00 

Zinc  acetate,  ~Zn(C2H802)2.7H20  N  154.7 

Salt  Solutions. 

Aluminum  chlorid,  A1C13.6H00  '      N/2  40.3  ' 

—nitrate,  A1(NO,),.9H80  N/2  62.5 

—potassium  sulfate,  KA1(S04)2.12H20  N/2  79 

$  Saturated  solutions  should  be  made  by  allowing  the  solvent  to  stand  in  contact 
with  the  solute  for  considerable  time,  the  mixture  being  shaken  occasionally.  After 
settling  thoroughly,  the  clear  solution  should  be  decanted  or  siphoned  off  into  another 
bottle  for  use. 


316  APPENDIX. 

Ammonium  chlorid,  NH4C1  N/2 

—nitrate,  NH4N03  N/2 

— sulfate,  (NH4)2S04  N/2 

Antimony  trichlorid,  SbCl3,  (see  page  318)  N/2 

— pentachlorid,  SbCl6,  (see  page  318)  N/2 

Arsenic  trichlorid,  AsCl3,  (see  page  318)  N/2 

— pentachlorid,  AsCl5,  (see  page  318)  N/2 

Barium  chlorid,  BaCl2.2H20  N/2 

—nitrate,  Ba(N03)2  N/2 

Bismuth  chlorid,  BiCl3,  (see  page  318)  N/2 

—nitrate,  Bi(N03)3.10H20,  (see  page  319)        N/2 

—sulfate,  Bi2(S04)3,  (see  page  319)  N/2 

Cadmium  chlorid,  CdCl2.2H20  N/2 

—nitrate,  Cd(N03)2.4H20  N/2 

—sulfate   (CdS04)3.8H20  N/2 

Calcium  chlorid,  CaCl2  N/2 

—nitrate,  Ca(N03)2  N/2 

Chromium  nitrate,  Cr(N03)3.9H20  N/2 

—sulfate,  Cr2(S04)3.6H20    '  N/2 

—potassium  sulfate,  KCr(SOJ2.12HaO  N/2 

Cobalt  acetate,  Co(C2H302)2.4H20  N/2 

—chlorid,  CoCl2.6H20  N/2 

—nitrate,  Co(N03)2.6H20  N/2 

—sulfate,  CoS04.7H20  N/2 

Copper  chlorid,  CuCl2.2H20  N/2 

—nitrate,  Cu(N03)2.6H20  N/2 

—sulfate,  CuS04.5H20*  N/2 

Ferric  chlorid,  FeCl3  N/2 

—nitrate,  Fe(N03)3.9H.,0  N/2 

—sulfate,  Fe2(SOJ3.xH20,  about  N/2 

Ferrous  chlorid,  FeCl2.4H20  N/2 

sulfate,  FeS04.7H20*  N/2 

Lead  acetate,  Pb(C2H302)2.3H20*  N/2 

—nitrate,  Pb(N03)2  N/2 

Lithium  chlorid,  LiCl  N/2 

—nitrate,  LiN08  N/2 

—sulfate,  LiS04  N/2 

Magnesium  chlorid,  MgCl2.6H20  N/2 

—nitrate,  Mg(N03)2.6H20  N/2 

—sulfate,  MgS04.7H20  ^       N/2 

Manganese  chlorid,  MnCl2.4H,0  N/2 

—nitrate,  Mn(N03)2.6H20  N/2 

—sulfate,  MnS04.4H20  N/2 

Mercuric  chlorid,  HgCl2,  about  N/2  sat. 

*  Should  contain  a  little  free  acid. 

$  Saturated  solutions  should  be  made  by  allowing  the  solvent  to  stand  in  contact 
with  the  solute  for  considerable  time,  the  mixture  being  shaken  occasionally.  After 
settling  thoroughly,  the  clear  solution  should  be  decanted  or  siphoned  off  into  another 
bottle  for  use. 


BEAGENTS   IN   SOLUTION.  317 

—nitrate,  (Hg(N08)2)2.H20*  N/2  83.3 

Mercurous  nitrate,  (HgN03)2.2H20,  (see  page  319)  N/2  135.5 

Nickel  chlorid,  NiCl2.6H20  N/2  59.5 

—nitrate,  Ni(N03)2.6H2O  N/2  72.7 

—  sulfate,  NiS04.7H20  N/2  70.2 
Potassium  arsenate,  K2HAs04  N/2  36.4 

—  arsenite  (meta),  KAs02  N/2  24.4 

—  bromid,  KBr  N/2  59.6 
—chlorate,  KC108,  about  N/2                            sat.  sol^ 
—chlorid,  KC1  N/2  37.3 

—  chromate,  K2Cr04  N/2  48.7 
Potassium  cyanate,  KOCN  N/2  40.6 

—  cyanid,  KCN  N/2  32.6 

—  dichromate,  K2Cr207  N/2  36.9 

—  ferricyanid,  K3(Fe  =  C6N6)  N/2  54.9 

—  ferrocyanid,  K4(Fe=  C0N6).3H20  N/2  52.9 

—  fluorid,  KF  N/2  29.1 

—  iodid,  KI  N/2  83.1 
—nitrate,  KN03  N/2  50.6 
—nitrite,  KN02  N/2  42.6 
—permanganate,  KMn04  N/2  39.6 
—sulfate,  K,S04  N/2  43.6 

—  sulfid,  K2S  N/2  27.6 

—  sulfite,  K2S03.2H20  N/2  48.6 

—  thiocyanate,  KSCN  N/2  48.6 
Silver  nitrate,  AgN03  N/2  84.9 
Sodium  acetate,  Na(C2H302).3H20  N/2  68 

—arsenate,  Na2HAs04.12H20  N/2  67 

—arsenite  (meta),  NaAsO»2  N/2  21.7 

—  borate  (tetra),  Na2B407.10H20  N/2  15.9 
—bromid,  NaBr  N/2  51.5 
—carbonate  (dry),  Na2CO3  N/2  26.5 
—carbonate   (crystals),  Na2CO,.10H,0  N/2  71.5 
—chlorid,  NaCl  N/2  29.3 
—fluorid,  NaF  N/2  21 

—  oxalate,  Na2C204  N/2  33.5 
—phosphate  (acid),  Na2HP04.12H2O  N/2  59.5 
—nitrate,  NaN08                                                  '  N/2  42.5 
—nitrite,  NaNO,  N/2  34.5 
—silicate,  Na2Si03                                      *           N/2  30.6 

—  silicofluorid,  Na2SiFa,  about  N/14  sat. 


*  Should  contain  some  free  acid. 

J  Saturated  solutions  should  be  made  by  allowing  the  solvent  to  stand  in  contact 
with  the  solute  for  considerable  time,  the  mixture  being  shaken  occasionally.  After 
settling  thoroughly,  the  clear  solution  should  be  decanted  or  siphoned  off  into  another 
bottle  for  use. 


318  APPENDIX. 

Sodium  sulfate,  Na2S04.10H20  N/2  80.5 

— sulfid,  Na,S  N/2  19.5 

— sulfite,  Na2S03.8H20  N/2  67.5 

— tartrate,  Na2(C4H406).2H20  N/2  57.5 

—potassium  tartrate,  NaK(C4H406).4H20          N/2  70.5 

— thiosulfate,  Na2S203.5H2O  N/2  62 

Stannic  chlorid,  SnCl4,  (see  page  319)  N/2  32.7 

Stannous  chlorid,  SnCl2.2H20,  (see  page  319)  N/2  56.5 

Strontium  chlorid,  SrCl2  N/2  39.7 

—nitrate,  Sr(N03)2.4H20  N/2  70.9 

Zinc  chlorid,  ZnCl2  N/2  34.1 

—nitrate,  Zn(NO,)2.6H20  N/2  74.4 

—sulfate,  ZnS04.7H20  N/2  71.8 

Reagents  Requiring  Special  Methods  of  Preparation. 

Ammonium  carbonate.  Dissolve  192  grams  of  the  solid  in  80  ccm.  of  cone. 
NH4OH  and  500  ccm.  of  water,  and,  after  solution  is  complete,  dilute  to  1  litre 
with  water. 

Ammonium  molybdate.  Dilute  150  ccm.  of  cone.  NH.OH  with  an  equal 
volume  of  water  and  dissolve  in  this  mixture  72  grams  of  molybdic  acid  anhy- 
drid,  Mo03.  Dilute  250  ccm.  of  cone.  HN03  with  twice  its  volume  of  water  and 
pour  the  first  solution  slowly  into  the  second  solution,  stirring  constantly.  Allow 
to  stand  in  a  warm  place  for  a  day  or  two,  decant  the  clear  liquid  above  any 
precipitate  that  may  have  formed  and  dilute  to  1  litre. 

Ammonium  sulfid.  Saturate  a  volume  of  cone.  NH4OH  with  H2S  by  passing 
the  gas  into  the  liquid  until  it  is  no  longer  absorbed.  Add  a  volume  of  cone. 
NH4OH  equal  to  the  resulting  solution  and  dilute  with  three  volumes  of  water. 

Ammonium  polysulfid.  "Yellow  ammonium  sulfid"  is  made  by  dissolving 
about  a  gram  of  finely  powdered  sulfur  in  each  litre  of  the  above  solution  before 
diluting  it.  After  the  sulfur  has  dissolved  it  may  be  diluted  with  an  equal  vol- 
ume of  water. 

Antimony  pentachlorid.  Dissolve  29.8  grams  of  SbCl5  in  400  ccm.  of  cone. 
HC1  mixed  with  an  equal  volume  of  water.  After  the  solution  is  complete, 
dilute  to  1  litre.  Or,  dissolve  32  grams  of  Sb205  in  400  ccm.  of  cone.  HC1  mixed 
with  400  ccm.  of  water  and  dilute  to  1  litre. 

Antimony  trichlorid.  Dissolve  37.6  grams  of  the  crystals  in  250  ccm.  of  cone. 
HC1  mixed  with  an  equal  volume  of  water.  After  solution  is  complete  dilute  to 
1  litre. 

Arsenic  pentachlorid  (?).  Dissolve  36.4  grams  of  K2HAs04  in  water,  add  50 
ccm.  of  cone.  HC1  and  dilute  to  1  litre. 

Arsenic  trichlorid.  Dissolve  33  grams  of  As203  in  50  ccm.  of  cone.  HC1 
diluted  with  10  times  its  volume  of  water.  When  the  oxid  has  entirely  dissolved, 
dilute  to  1  litre.  Or,  dissolve  24.4  grams  of  KAs02  in  water,  add  50  ccm.  of 
cone.  HC1  and  dilute  to  1  litre. 

Bismuth  chlorid.  Dissolve  52.5  grams  of  the  crystals  in  250  ccm.  of  water 
to  which  100  ccm.  of  cone.  HC1  has  been  added  and  dilute  to  1  litre. 


EEAGENTS   IN   SOLUTION.  319 

Bismuth  nitrate.  Dissolve  95.8  grams  of  Bi(N03)3.10H20  or  50.8  grams  of 
the  "  subnitrate,"  BiN03(OH)2,  in  250  ecm.  of  water  mixed  with  50  ccm.  of  cone. 
HN03.  When  solution  is  complete,  dilute  to  1  litre. 

Bismuth  sulfate.  Dissolve  58.7  grams  of  the  salt  in  250  ccm.  of  water  mixed 
with  25  ccm.  cone.  H2S04  and,  when  solution  is  complete,  dilute  to  1  litre. 

Brucin.    Dissolve  about  1  gram  of  the  substance  in  100  ccm.  of  cone.  H2S04. 

Devarda's  alloy.  Fuse  together  in  a  covered  earthenware  crucible  450 
grams  of  aluminum,  50  grams  of  zinc  and  500  grams  of  copper.  After  the  mix- 
ture has  become  liquid,  pour  out  upon  a  cold  surface  and  pulverize  when  cold. 

lodin  water.  Dissolve  50  grams  of  KI  in  a  litre  of  water  and  saturate  with 
iodin  by  allowing  the  solution  to  stand  in  contact  with  an  excess  of  iodin  crys- 
tals for  some  time  with  occasional  shaking.  Afterwards  decant  the  clear  dark- 
brown  solution. 

"  Magnesia  mixture."  Dissolve  101.8  grams  of  MgCl2.6H20  and  107  grams 
NH4C1  in  water,  add  400  ccm.  cone.  NH4OH  and  dilute  to  1  litre. 

Manganese  chlorid  (concentrated).  Saturate  cone.  HC1  with  the  crystallized 
salt,  MnCl2.4H20. 

Mercurous  nitrate.  Dissolve  135.5  grams  of  the  crystallized  salt  in  a  half 
litre  of  water  to  which  25  ccm.  of  cone.  HN03  has  been  added.  Dilute  to  1 
litre  and  add  50-100  grams  of  mercury.  Or,  treat  100  grams  of  mercury  with  25 
ccm.  of  cone.  HN03  diluted  to  125  ccm.  with  water  and  maintain  near  the  boil- 
ing temperature  until  the  metal  has  disappeared,  adding  more  acid  in  small 
amounts  if  necessary.  Decant  the  clear  liquid  and  add  50-100  grams  of  mer- 
cury and  dilute  to  1  litre. 

Phenylenediamin  (meta).  Dissolve  2-3  grams  of  the  solid  in  10  ccm.  of  cone. 
HC1  and  90  ccm.  of  water. 

Potassioscope  Liquid.  Dissolve  about  0.5  gram  of  Gentian  Blue  in  a  litre  of 
95  per  cent,  alcohol.  Place  in  well-corked  flat  bottles  of  such  size  that  a  layer  of 
solution  about  1  inch  thick  may  be  obtained.  Six  ounce  vials  are  satisfactory. 
Bottles  with  curved  surfaces  should  be  avoided. 

A  solution  of  the  dye  in  dilute  alcohol  will  precipitate  on  standing. 

Stannous  chlorid  (concentrated  for  Bettendorff  test  for  arsenious  acid). 
Treat  600  grams  of  granulated  tin  with  a  litre  of  cone.  HC1  and  maintain  the 
mixture  at  a  temperature  sufficiently  high  to  keep  up  a  vigorous  action  of  the 
acid  on  the  metal.  When  hydrogen  is  no  longer  liberated,  decant  the  solution 
and  add  a  few  ccm.  of  cone.  HC1  and  a  few  pieces  of  the  metal  to  the  bottle  in 
which  it  is  to  be  kept.  Or,  if  only  a  small  amount  of  the  reagent  is  desired,  dis- 
solve 113  grams  of  SnCl2.2H20  in  75  ccm.  of  cone.  HC1  and  add  a  few  pieces  of 
the  metal. 

Stannous  chlorid  (general  reagent).  Dilute  50  ccm.  of  the  previous  solution 
with  an  equal  volume  of  cone.  HC1  and  dilute  to  1  litre  with  water,  keeping  a 
few  pieces  of  metallic  tin  in  contact  with  the  solution. 

Stannic  chlorid.  Add  bromin  water  to  50  ccm.  of  the  concentrated  solution 
of  SnCl2  prepared  as  above  until  a  permanent  yellow  color  is  produced.  Boil 
off  the  excess  of  bromin  (indicated  by  disappearance  of  color)  and  dilute  to  1 
litre. 

Starch  solution.  Finely  pulverize  1  gram  of  starch,  rub  up  to  a  thin  mixture 
with  about  10  ccm.  water  and  pour  immediately  into  100  ccni.  of  boiling  water. 
Boil  for  5-10  minutes. 


320  APPENDIX. 

Starch-iodid  mixture.  Add  1  gram  of  KI  to  100  ccm.  of  the  starch  solution 
prepared  as  previously  directed. 

Sulfanilic  acid  mixture.  Dissolve  0.5  gram  of  sulfanilic  acid  in  150  ccm.  of 
dilute  acetic  acid  made  by  mixing  50  ccm.  of  the  80  per  cent,  acid  with  100  ccm. 
of  water.  Boil  0.1  gram  of  a-naphthylamin  with  25  ccm.  of  water  and  decant 
the  colorless  solution  from  the  violet  residue  and  add  50  ccm.  of  dilute  acetic 
acid  of  the  strength  given  to  the  solution. 

The  two  solutions  are  then  mixed  and  should  be  nearly  colorless.  If  a  pink 
tinge  should  appear  on  standing,  it  should  be  removed  by  shaking  the  solution 
with  zinc  dust  and  then  filtering.  Use  1  ccm.  in  testing  for  nitrous  acid. 

Sulfuric  acid.  Since  a  large  amount  of  heat  is  produced  when  cone.  H2S04  is 
diluted,  to  avoid  dangerous  spattering  and  breaking  of  containers,  the  cone,  acid 
should  always  be  added  to  the  necessary  volume  of  water  slowly  and  with  con- 
stant stirring. 


SOLID    REAGENTS.  321 

REAGENTS   IN   SOLID   FORM. 

Ammonium  chlcrid,  NH4C1. 

—nitrate,  NH4N03. 
Barium  hydroxid,  Ba(OH)2.8H20. 
Borax,  Na2B407.10H20. 
Calcium  carbonate,  CaC03. 

— fluorid,  CaF2  (fluorspar). 

—hydroxid,  Ca(OH)2,  "slacked  lime." 

— oxid,  CaO,  lime. 
Copper,  foil. 

— oxid,  CuO. 

Devarda's  alloy  (see  page  319). 
Ferrous  sulfate,  FeS04.7H20. 
Iron  filings. 

— nails. 

Lead  dioxid,  Pb02. 
Lime  (see  calcium  oxid). 
Litmus  paper. 

Microcosmic  salt,  HNa(NH4)P04.4H20. 
Paraffine. 
Potassium  carbonate,  K2C03. 

—chlorate,  KC103. 

— cyanid,  KCN. 

— dichromate,  K2Crs07. 

— ferricyanid,  K^Fe^C.NJ. 

—nitrate,  KN03. 

—sulfate  (acid),  KHS04. 
Sodium,  metal. 

— carbonate,  Na,C03. 

—carbonate  (acid),  NaHC03. 

— and  potassium  carbonate  mixture  for  fusions. 

—hydroxid,  NaOH. 

— meta-phosphate,  NaP03  (prepared  from  microcosmic  salt)  (see  page  78). 

— peroxid,  Na202. 

— sulfite,  Na0S03.8H20. 
Sulfur,  powdered. 
Tumeric  paper. 
Zinc,  dust. 

— metal. 


22 


322 


APPENDIX. 


SPECIFIC    GRAVITY    OF    SOLUTIONS   AT    15°  C. 

SULFURIC  ACID.     (Lunge  and  Isler.} 


Specific 
Gravity. 

Per  Cent. 
H2S04. 

Specific 
Gravity. 

Per  Cent.       1 
H2S04. 

1       Specific 
Gravity. 

Per  Cent. 
H2S04. 

1.000 

0.09 

.320 

41.50 

1.640 

71.99 

1.010 

1.57 

.330 

42.66 

1.650 

72.88 

1.020 

3.03 

.340 

43.74 

1.660 

73.64 

1.030 

4.49 

.350 

44.82 

1.670 

74.51 

1.040 

5.96 

.360 

45.88 

1.680 

75.42 

1.050 

7.37 

.370 

46.94 

1.690 

76.30 

1.060 

8.77 

.380 

48.00 

•  1.700 

77.17 

1.070 

10.19 

.390 

49.06 

1.710 

78.04 

1.080 

11.60 

1.400 

50.11 

1.720 

•  78.92 

1.090 

12.99 

1.410 

51.15 

1.730 

79.80 

1.100 

14.35 

1.420 

52.15 

1.740 

80.68 

1.110 

15.71 

1.430 

53.11 

1.750 

81.56 

1.120 

17.01 

1.440 

54.07 

1.760 

82.44 

1.130 

18.31 

1.450 

55.03 

1.770 

83.32 

1.140 

19.61 

1.460 

55.97 

1.780 

84.50 

1.150 

20.91 

1.470 

56.90 

1.790 

85.70 

1.160 

22.19 

1.480 

57.83 

1.800 

86.90 

1.170 

23.47 

1.490 

58.74 

1.810 

88.30 

1.180 

24.76 

1.500 

59.70 

1.820 

90.05 

1.190 

26.04 

1.510 

60.65 

1.825 

91.00 

1.200 

27.32 

1.520 

61.59 

1.830 

92.10 

1.210 

28.58 

1.530 

62.53 

1.834 

93,05 

1.220 

29.84 

1.540 

63.43 

1.837 

94.20 

1.230 

31.11 

1.550 

64.26 

1.839 

95.00 

1.240 

32.28 

1.560 

65.08 

1.840 

95.60 

1.250 

33.40 

1.570 

65.90 

1.841 

97.00 

1.260 

34.57 

1.580 

66.71 

1.8415 

98.20 

1.270 

35.71 

1.590 

67.59 

1.841 

98.70 

1.280 

36.87 

1.600 

68.51 

1.840 

99.20 

1.290 

38.03 

1.610 

69.43 

1.839 

99.70 

1.300 

39.19 

1.620    ' 

70.32 

1.310 

40.35 

1.630 

71.16 

HYDROCHLORIC  ACID.      (Lunge  and  Marchlewski.) 


Specific 
Gravity. 

Per  Cent. 
HC1. 

Specific 
Gravity. 

Per  Cent. 
HC1. 

Specific 
Gravity. 

Per  Cent. 
HC1. 

1.000 

0.16 

1.070 

14.17 

1.140 

27.66 

1.005 

1.15 

1.075 

15.16 

1.145 

28.61 

.010 

2.14 

1.080 

16.15 

1.150 

29.57 

.015 

3.12 

1.085 

17.13 

1.155 

30.55 

.020 

4.13 

1.090 

18.11 

1.160 

31.52 

.025 

5.15 

1.095 

19.06 

1,165 

32.49 

.030 

6.15 

1.100 

20.01 

1.170 

33.46 

.035 

7.15 

1.105 

20.97 

1.175 

34.42 

.040 

8.16 

1.110 

21.92 

1.180 

35.39 

.045 

9.16 

1.115 

22.86 

1.185 

36.31 

.050 

10.17 

1.120 

23.82 

1.190 

37.23 

.055 

11.18 

1.125 

24.78 

1.195 

38.16 

.060 

12.19 

1.130 

25.75 

1.200 

39.11 

1.065 

13.19 

1.135 

26.70 

SPECIFIC  GKAVJTY  OF  SOLUTIONS. 


323 


NITRIC  ACID.     (Lunge  and  Eey.) 


Specific 
Gravity. 

Per  Cent. 
HNO8. 

Specific 
Gravity. 

Per  Cent. 
HN03. 

Specific 
Gravity. 

Per  Cent. 
HN03. 

1.00 

0.00 

1.18 

29.38 

1.36 

57.57 

1.01 

1.90 

1.19 

30.88 

1.37 

59.39 

1.02 

3.70 

1.20 

32.36 

1.38 

61.27 

1.03 

5.50 

1.21 

33.82 

1.39 

63.23 

1.04 

7.26 

1.22 

35.28 

1.40 

65.30 

1.05 

8.99 

1.23 

36.78 

1.41 

67.50 

1.06 

10.68 

1.24 

38.29 

1.42 

69.80 

1.07 

12.33 

1.25 

39.82 

1.43 

72.17 

1.08 

13.95 

1.26 

41.34 

1.44 

74.68 

1.09 

15.53 

1.27 

42.87 

1.45 

77..2S 

1.10 

17.11 

.1.28 

44.41 

1.46 

79.98 

1.11 

18.67 

1.29 

45.95 

1.47 

82.90 

.12 

20.23 

1.30 

47.49 

1.48 

86.05 

.13 

21.77 

1.31 

49.07 

1.49 

89.60 

.14 

23.31 

1.32 

50.71 

1.50 

94.09 

.15 

24.84 

1.33 

52.37 

1.51 

98.10 

.16 

26.36 

1.34 

54.07 

1.52 

99.67 

1.17 

27.88 

1.35 

55.79 

AMMONIA  SOLUTIONS.     (Lunge  and  Wiernick^ 


Specific 
Gievity. 

Per  Cent. 
NH3. 

Specific 
Gravity. 

Per  Cent. 
KH3. 

Specific 
Gravity. 

Per  Cent. 
NH3. 

0.882 

34.95 

0.922 

21.12 

0.962 

*9.35 

0.884 

34.10 

0.924                  20.49 

0.964 

8.84 

0.886 

33.25 

0.926 

19.87 

0.966 

8.33 

0.888 

32.50 

0.928 

19.25 

0.968 

7.82 

0.890 

31.75 

0.930 

18.64 

0.970 

7.31 

0.892 

31.05 

0.932 

18.03 

0.972 

6.80 

0.894 

30.37 

0.934 

17.42 

0.974 

6.30 

0.896 

29.69 

0.936 

16.82 

0.976 

5.80 

0.898 

29.01 

0.938 

16.22 

0.978 

5.30 

0.900 

28.33 

0.940 

15.63 

0.980 

4.80 

0.902 

27.65 

0.942 

15.04 

0.982 

4.30 

0.904 

26.98 

0.944 

14.46 

0.984 

3.80 

0.906 

26.31 

0.946 

13.88 

0.986 

3.30 

0.908 

25.65 

0.948 

13.31 

0.988 

2.80 

0.910 

24.99 

0.950 

12.74 

0.990 

2.31 

0.912 

24.33 

0.952 

12.17 

0.992 

1.84 

0.914 

23.68 

0.954 

11.60 

0.994 

1.37 

0.916 

23.03 

0.956 

11.03 

0.996 

0.91 

0.918 

22.39 

0.958 

10.47 

0.998 

0.45 

0.920 

21.75 

0.960 

9.91 

1.000 

0.00 

324 


APPENDIX. 


LIST  OF  ELEMENTS.*     1906. 


Name. 

Sym- 
bol. 

Valences. 

At.Wt. 

Name. 

Sym- 
bol. 

Valences. 

At.Wt. 

Aluminum  

Al 

3 

27  1 

Neon 

Ne 

o 

20 

Antimony  

Sb 

3-5 

120  2 

Xickel 

Ni 

2-3 

58  7 

Argon 

A 

o 

39  9 

Nh 

5  ^ 

Q4 

Arsenic 

As 

3-5 

75 

Nitrogen 

N 

5  H 

Barium  

Ba 

2 

137  4 

Osmium 

Os 

864 

m" 

Beryllium  

Be 

2 

9.1 

Oxygen  

o 

2 

16 

Bismuth  

Bi 

3 

208  5 

Palladium 

Pd 

4  2 

106  5 

Boron  

B 

3 

11 

Phosphorus 

p 

5  3 

31 

Bromin  

Br 

1-5 

80 

Platinum 

Pt 

4-2 

194  8 

Cadmium  

Cd 

2 

112  4 

Potassium 

K 

1 

39  2 

Caesium  

Cs 

1 

132  9 

Praseodymium 

Pr 

3 

140  5 

Calcium 

Ca 

2 

40  1 

Hadium 

Ra 

2 

225 

Carbon 

c 

4-2 

12 

Rhodium 

Rh 

3 

103 

Cerium    

Ce 

4-3 

140  2 

Rubidium 

Rb 

1 

85  5 

(  hlorin  

Cl 

1-5-7-3 

35  5 

Ruthenium 

Ru 

4  3-6-8 

101  7 

Chromium   ... 

Cr 

3_6-7-2 

52  1 

Samarium 

Sa 

3 

150  3 

Cobalt  

Co 

2-3 

59 

Scandium 

Sc 

3 

44  1 

Copper  

Cu 

2-1 

63.6 

Selenium 

Se 

2-6-4 

79  2 

Er 

3 

166 

Silicon  

Si 

4 

28  4 

Fluorin  

F 

1 

19 

SilTer.  

Aff 

1 

107  9 

Gadolinium  

Gd 

3 

156 

Sodium  

Na 

1 

23 

(  »iil  Hum 

Ga 

3 

70 

Strontium 

Sr 

2 

87  6 

Germanium  

Ge 

4-2 

72.5 

Sulfur  

s 

2-4-6 

32  1 

Gold 

Au 

1-3 

197  2 

Tantalum 

Ta 

5 

183 

Helium  .. 

He 

o 

4 

Tellurium 

Te 

2-4-6 

127  6 

Hydrogen 

H 

I 

1 

Terbium 

Tb 

3 

160 

Indium  

In 

3 

115 

Thallium 

Tl1 

1-3 

204  1 

lodin  

I 

1-5-7 

127 

Thorium  

Th 

4-3 

232.5 

Iridium  

Ir 

3-4-2 

193 

Thullium  

Tu 

3 

171 

Iron  

Fe 

3-2 

55.9 

Tin   

Sn 

2-4 

119 

Krypton 

Kr 

o 

81  8 

Titanium 

Ti 

4 

48  1 

Lanthanum 

La 

3 

138  9 

Tungsten 

W 

642 

184 

Lead 

Pb 

2  4 

206  9 

Uranium 

Tj 

6-4 

238  5 

Lithium 

Li 

1 

7 

Vanadium  v 

V 

5-3 

51  2 

Magnesium 

Me 

2 

24  4 

Xenon 

Xe 

o 

128 

Manganese  . 

Mn 

2-7-4-6-3 

55 

Ytterbium 

Yb 

3 

173 

Mercury  

Hg 

2-1 

200 

Yttrium 

Y 

3 

89 

Molybdenum  

Mo 

6-4-2 

96 

Zinc  

Zn 

2 

65  4 

Neodymium  

Nd 

3 

143.6 

Zirconium  

Zr 

4 

90.6 

NOMENCLATURE. 


The  elements  are  named  quite  arbitrarily,  except  that  the  ending  "  ium  "  or 
"  um  "  is  used  to  designated  the  possession  of  metallic  properties.  All  of  the 
more  recently-named  metals  are  so  indicated.  Very  often  an  element  is  named 
because  of  some  very  characteristic  property,  the  name  being  derived  usually 
from  Greek  words  signifying  the  possession  of  this  quality.  Chlorin  and  iodin 
come  from  the  Greek  words  meaning  "  green "  and  "  violet,"  both  of  these 
*  Calculated  from  the  basis,  O  =  16. 


NOMENCLATUEE.  325 

names  signifying  the  color  of  the  vapor  of  these  elements.  Bromin  is  from  the 
word  meaning  "  stench,"  referring  to  the  manner  in  which  this  substance  affects 
the  membranes  of  the  nose  and  throat.  Hydrogen,  oxygen  and  nitrogen  are 
compounded  from  the  verb  gennao,  meaning  "  to  produce,"  and  from  the  nouns 
meaning  "  water,"  "  acid,"  and  "  nitre,"  respectively.  Hydrogen  is  the  element 
which  forms  water  when  it  burns;  oxygen  is  a  constituent  of  very  many  acids, 
once  thought  to  be  present  in  all,  and  nitrogen  a  principal  component  of  nitre. 
Potassium,  sodium  and  aluminum  are  the  metals  present  in  potash,  soda  and 
alum.  Phosphorus  is  compounded  from  Greek  words,  and  means  "  the  light- 
bearer,"  referring  to  the  phosphorescence  of  the  element.  Chromium  is  from  the 
word  meaning  color,  suggesting  the  fact  that  most  compounds  of  this  element 
are  highly  colored.  Certain  of  the  elements  are  named  from  the  countries  in 
which  they  were  discovered,  as  magnesium  from  Magnesia;  scandium  from 
Scandanavia;  gallium  from  Gaul,  the  older  name  for  France;  germanium  from 
Germany.  Tellurium  comes  from  the  Greek  word  meaning  the  earth,  selenium, 
similarly,  signifies  the  moon,  and  helium  the  sun,  the  latter  name  being  given 
because  of  the  fact  that  this  element  was  known  to  exist  in  the  sun  before  it  was 
discovered  on  earth.  Gold,  silver,  copper,  lead,  mercury,  tin  and  iron  were 
known  to  the  ancients,  and  in  their  chemical  symbols  the  older  names  are  still 
retained. 

In  the  naming  of  chemical  compounds,  however,  certain  definite  rules  are  fol- 
lowed. All  binary  compounds  (£.  e.,  consisting  of  two  substances)  are  designated 
by  naming  the  metallic  element  first  followed  by  the  stem  of  the  non-metallic 
element  to  which  the  suffix,  "id,"  is  added,  as  HC1,  hydrogen  chlorid;  ZnS,  zinc 
sulfid;  CaO,  calcium  oxid;  Na(OH),  sodium  hydroxid,  the  hydroxyl  radicle  be- 
ing considered  as  a  unit. 

If  two  bodies  are  composed  of  the  same  elements  in  different  proportions, 
the  one  containing  the  greater  proportion  of  the  non-metallic  constituent  is 
designated  by  the  ending  "  ic,"  while  the  compound  with  the  smaller  proportion 
of  the  non-metallic  constituent  is  indicated  by  the  suffix  "  ous."  Thus  FeCls  is 
ferric  chlorid;  FeCl2,  ferrous  chlorid;  Sn02,  stannic  oxid;  SnO,  stannous  oxid; 
H2S04,  sulfuric  acid;  H2S03,  sulfurous  acid. 

When  more  than  two  compounds  containing  the  same  elements  exist,  they 
are  designated  by  the  prefixes,  "  per,"  meaning  "  higher,"  and  "  hypo,"  mean- 
ing "  lower." 

Salts  of  acids  ending  in  "ous"  are  indicated  by  the  ending  "ite"  while 
salts  of  "  ic  "  acids  end  in  "  ate." 

The  following  table  will  make  plain  the  usage: 

Formula.  Acid.                                                Salt.  Formula. 

HC1O  Hypochlorous  acid  Sodium  hypochlorite  NaCIO 

HC1O,  Chlorous  acid  Sodium           chlorite  NaC102 

HClo]  Chloric  acid  Sodium          chlorate  NaCIO, 

HC104  Perchloric  acid  Sodium    perchlorate  NaC104 


326 


APPENDIX. 


PHYSICAL   PROPERTIES    OF    THE    ELEMENTS. 


Element. 

Specific 
Gravity 

Melting 
Point. 

Boiling 
Point. 

Element. 

Specific 
Gravity. 

Melting 
Point. 

Boiling 
Point. 

Aluminum     

2  67 

657 

1470-1700 

Neon 

Antimony  

6  72 

432 

1500-1700 

Nickel  .    . 

8  85 

1484 

Arson 

187.9 

—  186  1 

Niobium  

7  06 

1950 

5  70 

500 

449  5 

Nitrogen 

214 

Btiri  Hill 

850 

950 

Osmium       ' 

22  48 

2500 

Beryllium 

1  85 

>  960 

Oxygen 

^>     230 

184 

Bismuth          .... 

9  65 

269 

1435 

Palladium 

11  40 

1586 

Boron 

2  58 

Infusible 

3500 

c 

1  82Y 

Bromin        

3  10 

—  7  3 

63 

Phosphorus...  < 

2  llK 

44.2 

2873 

Cadmium 

8  72 

321  7 

778 

Platinum     .    . 

21  48 

1770 

Capsium    

1  88 

26  37 

670 

Potassium   .    . 

0  87 

62  5 

667 

(Calcium  

780 

Praseodymium. 

940 

3  55c 

Radium  

2  25G 

3500 

Rhodium  

11.50 

2000 

1  70A 

Rubidium   

1.52 

38  5 

696 

Cerium  

6  65 

623 

Ruthenium  

11.20 

1900 

Chlorin  

—  102 

—  33  6 

Samarium  

Chromium 

6  47 

1515 

Scandium 

Cobalt 

8  55 

1530 

Selenium  .. 

4  48 

175 

690 

8.93 

1045 

2100 

Silicon  

2.49 

1200 

Erbium  

Silver  

10  50 

955 

Fluorin  

—  223 

•  —  187 

Sodium  

0  97 

95  6 

742 

Gadolinium.     ... 

Strontium  

2  55 

900 

Gallium  

5.95 

30.15 

Qnlfn-r                           J 

2.06o 

A  A  A     f> 

Germanium  
Gold  

5.47 
19.32 

900 
1065 

1350 

buiiur  < 
Tantalum  

1.96  M 
10.70 

iio—  iiy 
2250 

Helium 

*>  —  271 

^>     270 

Tellurium 

6  24 

446 

1390 

Hydrogen 

—  256  5 

252  5 

Terbium  

Indium 

7  40 

155 

800 

Thallium  

11.85 

302 

1600-180C 

lodin 

4.95 

1142 

184  4 

Thorium  

11  00 

Iridium 

21  83 

1950 

Thullium  

Iron     

7  80 

1804 

Tin   

7.29 

232 

1450-1  60C 

Krypton  

—169 

—151.7 

Titanium  

3000 

Lanthanum  
Lead 

6.10 
11  37 

810 
327 

1400-1600 

Tungsten  
Uranium       

17.20 

18  68 

1700 
800 

Lithium 

0  59 

186 

<  1400 

Vanadium  

5  50 

1630 

Magnesium 

1  75 

632.6 

1100 

Xenon  

—140 

—109.1 

7  25 

1900 

Ytterbium  . 

13  57 

38  85 

357 

Yttrium 

8  65 

Zinc       

715 

419 

918 

Neodymium  

840 

Zirconium  

4.15 

1500 

POTENTIAL    SEKIES.  327 


POTENTIAL   SERIES. 

The  following  list  shows  the  comparative  readiness  with  which  various  metals 
tend  to  take  on  charges  of  plus  electricity  and  pass  into  the  ionic  condition,  the 
element  possessing  this  property  in  the  maximum  degree  standing  first.  Every 
metal  in  this  list  is  said  to  be  positive  to  all  the  metals  below  it,  i.  e.}  tends  to 
pass  more  readily  into  the  ionic  condition,  or  liberates  more  energy  in  taking  on 
the  electrical  charge.  Since  this  determines  the  potential  when  two  metals  are 
joined  together  in  a  cell,  this  arrangement  is  known  as  the  Potential  Series. 

Because  each  element  tends  to  pass  more  readily  into  the  ionic  condition  than 
any  element  standing  lower  in  the  series,  whenever  a  metal  is  introduced  into  a 
solution  of  a  salt  of  another  metal,  theoretically,  the  more  positive  metal  should 
replace  the  less  positive,  the  former  should  take  the  charge  from  the  latter, 
the  more  positive  dissolving  as  the  less  positive  is  precipitated  from  solution. 
Under  appropriate  condition  this  can  generally  be  accomplished  in  most  eases 
but  the  action  is  most  vigorous  between  metals  widely  separated  in  the  Potential 
Series.  For  the  same  reason,  those  metals  standing  higher  than  hydrogen  dis- 
solve readily  in  acids  unless  an  insoluble  product  forming  a  protective  coating 
about  the  metal  is  produced  ly  the  reaction.  On  the  other  hand,  those  metals 
standing  lower  than  hydrogen  in  this  series  are  not  readily  attacked  by  acids 
when  the  action  is  due  to  hydrogen  ions  only.  (See  footnote,  page  17.) 

1  Potassium  10  Zinc  18  Arsenic 

2  Sodium  11  Cadmium  19  Antimony 

3  Lithium  12  Iron  20  Tin 

4  Barium  13  Cobalt  21  Copper 

5  Strontium  14  Nickel  22  Mercury 

6  Calcium  15  Lead  23  Silver 

7  Magnesium  16  Hydrogen  24  Platinum 

8  Aluminum  17  Bismuth  25  Gold 

9  Manganese 

Few  non-metallic  elements  form  simple  anions.  The  following  order* has  been 
determined,  however,  among  those  that  are  well  known: 

1  Fluorin  3  Bromin  5  lodin 

2  Chlorin  4  Oxygen  6  Sulfur 


328 


APPENDIX. 


METRIC   MEASURES   WITH   ENGLISH   EQUIVALENTS. 

Measures  of  Length. 

1  Millimeter,  mm.  =     .03937  inches. 

1  Centimeter,  cm.  =  10  mm.  =  .3937  inches. 
1  Decimeter,  dm.  =10  cm.  =  3.9371  inches. 
1  Meter,  m.  =10  dm.  =39.3708  inches. 

1  Kilometer,   km.   =1000m.=     .6214  miles. 


Measures  of  Volume. 

1  Cubic  centimeter,  ccm.  .06103  cubic  inches. 

1  Cubic  deleter,  (liter),  1.  =1000  ccm.  =  {  OJJ -J- £j  or 

1  Cubic  meter                          =10001.       =  35.3166  cubic  feet. 


Measures  of  Weight. 

1  Milligram,  mg.  =     .0154  grains. 

1  Gram  =  1000  mg.       =  15.432  grains. 

1  Kilogram   (kilo)   =1000  grams.  =   2.2046  pounds  av. 


ENGLISH   MEASURES   WITH   METRIC    EQUIVALENTS. 

1  Inch  =  25.399  millimeters.  1  Cubic  inch  =  16.386  cubic  centimeters. 

1  Foot  =      .3048  meters.  1  Cubic  foot  =  28.315  liters. 

1  Mile  =   1.609  kilometers.  1  Quart          =      .9463  liters. 

1  Grain  =       .0648  grams. 

1  Ounce  av.  =  28.3496  grams. 

1  Pound    av.  =       .4536  kilograms. 


Apothecaries'  Weights  and  Measures  with  Metric  Equivalents. 


1  Grain  =       .0648  grams. 

1  Scruple        =     1.296  grams. 
1  Drachm        =     3.888  grams. 
1  Ounce  troy  =   31.1035  grams. 
1  Pound  troy  =  373.2418  grams. 


1  Minim  =        .0616  cubic  centimeters. 

1  Fluid   drachm  =      3.6965  cubic  centimeters. 

1  Fluid  ounce  =   29.572  cubic  centimeters. 

1  Pint  =  473.11  cubic  centimeters. 


FACTOES. 


329 


CONVERSION   FACTORS. 


TO   CHANGE 

TO 

MULTIPLY   BY 

Inches 

Centimeters 

2.54 

Feet 

Meters 

.305 

Miles 

Kilometers 

1.609 

Meters 

Inches 

39.37 

Kilometers 

Miles 

.621 

Square  inches 

Square  centimeters 

6.452 

Square  yards 

Square  meters 

.836 

Square  centimeters 

Square  inches 

.155 

Square  meters 

Square  yards 

1.196 

Cubic  inches 

Cubic  centimeters 

16.386 

Cubic  yards 

Cubic  meters 

.765 

Cubic  centimeters 

Cubic  inches 

.061 

Cubic  meters 

Cubic  yards 

1.308 

Fluid  ounces 

Cubic  centimeters 

29.57 

8uarts 

Liters 

.946 

ubic  centimeters 

Fluid  ounces 

.034 

Liters 

Quarts 

1.057 

Grains 

Milligrams 

64.799 

Ounces  (av.) 

Grams 

28.35 

Pounds   (av.) 

Kilograms 

.454 

Ounces  (troy) 

Grams 

31.103 

Pounds  (troy) 

Kilograms 

.373 

Grams 

Grains 

15.432 

Kilograms 

Pounds 

2.205 

330 


APPENDIX. 


TABLE    OF    SOLUBILITIES. 


1 

1 

K 
0 

o 
II 

CO 

fc 

0 

i 

p 

5 

7 

g 

| 

1 

? 

(0 

5 

O 

p" 

o 
o 

0 

o 

1 

1 

1 

1 

1 

1 

1 

1 

1 

Na  — 

1 

1 

1 

1 

1 

1 

i 

1 

-1 

1 

1 

i 

1 

I 

1 

1 

K- 

1 

1 

1 

1 

1 

1 

i 

1 

1 

1 

1 

i 

1 

1 

1 

1 

NEL  

1 

1 

1 

1 

1 

1 

i 

1 

1 

1 

1 

i 

1 

1 

1 

1 

Mg  = 

3 

3 

1 

4 

1 

1 

i 

1 

1 

1 

1 

i 

1 

1 

1 

1 

Ca  = 

2 

1 

2 

4 

1 

i 

1 

1 

1 

1 

i 

1 

1 

1 

1 

Sr  = 

2 

1 

1-2 

4 

1 

i 

1 

1 

1 

1 

i 

1 

1 

J 

1 

Ba  = 

1 

1 

1 

4 

1 

i 

2 

1 

1 

1 

i 

1 

1 

1 

1 

Zn  = 

3 

3 

3 

3 

1 

i 

3 

3 

1 

3 

3 

— 

1 

— 

1 

Mn  = 

3 

3 

3 

3 

1 

i 

3 

3 

1 

4 

3 



1 



1 

Ni  = 

3 

3 

4 

2 

1 

i 

4 

3 

1 

4 

4 

— 

1 

— 

1 

Co  = 

3 

3 

4 

2 

1 

i 

4 

3 

1 

4 

4 

— 

1 

— 

1 

Fe  = 

3 

3 

3 

2 

1 

i 

4 

3 

1 

4 

4 

— 

1 

— 

1 

Fe  = 

3 

4 



1 

1 

i 

a 

a 

1 

1 

4 



1 



1 

Cr  = 

3 

4 



1 

1 

i 

a 

a 

1 

— 

— 

— 

1 

— 

1 

Al  == 

3 

4 



1 

1 

i 

a 

a 

1 

— 

a 

— 

I 

— 

1 

Sn  = 

3 

3 

3 

1 

1 

2 

a 

a 

— 

3 

3 

— 

— 

— 

— 

Sn^ 

3 

4 

3 

1 

1 

1 

1 

a 

a 



3 

3 





— 

— 

SI)  EEE 

3 

3 

3 

1 

J6 

1 

& 

16 



— 

— 

— 

— 

— 

— 

— 

— 

As  = 

2 

2 

4 

1 

1 

1 

1 

— 

— 

— 

— 

— 

— 

— 

— 

— 

Cd  = 

3 

3 

3 

2 

1 

1 

1 

3 

3 

1 

3 

3 

— 

1 

— 

1 

Cu  = 

3 

3 

3 

3 

1 

1 

4 

4 

3 

3 

4 

4 



1 



1 

Bi  = 

3 

3 

3 

1 

16 

1 

1 

a 

1 

3 

3 



1 

— 

1 

Pb  = 

3 

3 

3 

3 

2-4 

4 

2-4 

3 

3 

1 

3 

3 

— 

1 

— 

1 

Hg  = 

3 

4 

1 

1 

1 

2 

1 

1 

2-3 

3 

1 

— 

1 

— 

1 

Hg- 





— 

3 

4 

4 

2 

— 

3 

3 

4 

4 

— 

1 

— 

— 

Ag- 

— 

3 

o 

1 

5 

5 

5 

5 

3 

4 

3 

4 

2 

1 

1 

1 

—  =  Unknown,  rare,  or  not  readily  formed.  1  =  Eeadily  soluble  in  water. 
2  =  Sparingly  soluble  in  water ;  generally  dissolve  readily  in  very  dilute  acids.  3  = 
Practically  insoluble  in  water,  but  dissolve  readily  in  4N  mineral  acids.  4  —  Eequire 
hot,  strong  acids  for  appreciable  solvent  effect.  5  =  Practically  insoluble  in  acids. 

a  •=  Hydroxid  is  precipitated.     &  =  Insoluble  basic  salt  likely  to  form. 


SOLUBILITIES. 


331 


TABLE    OF    SOLUBILITIES. 


V' 

• 

0* 

o 
5 
1 

o 

a 

1 

i 

q 

II 

i 

| 

o" 
33 
II 

0° 

o 
II 

q 
o 

^ 

o 

II 

W 

cT 
1 

2 

I 

1 

0 

w 

33 
II 

1 

1 

i 

i 

1 

l 

1 

1 

1 

1 

i 

1 

1 

1 

1 

-Na 

1 

1 

i 

i 

1 

1 

1 

1 

1 

1 

i 

1 

1 

1 

2-4 

K 

1 

1 

i 

i 

1 

— 

1 

1 

1 

1 

i 

1 

1 

1 

1 

-NH4 

1 

1 

2 

i 

1 

3 

3 

2 

3 

1 

3 

3 

3 

2 

1 

=  Mg 

1 

1 

3 

2-4 

1 

3 

3 

3 

3 

1 

3 

3 

3 

2 

1 

=  Ca 

1 

1 

3 

5 

2 

3 

3 

3 

3 

1 

3 

3 

3 

2 

1 

=  Sr 

1 

1 

3 

2 

5 

3 

3 

3 

3 

3 

1 

3 

3 

3 

2 

2-4 

=  Ba 

1 

1 

2 

1 

1 

1 

3 

3 

3 

3 

1 

3 

3 

3 

2 

1 

=  Zn 

1 

1 

2 

1 

1 

1 

3 

3 

2 

3 

1 

3 

3 

3 

2 

1 

=  Mn 

1 

1 

3 

1 

1 

3 

3 

3 

3 

3 

1 

3 

3 

3 

3 

1 

—  Ni 

1 

1 

3 

1 

1 

3 

3 

3 

3 

3 

1 

3 

3 

3 

3 

1 

=  Co 

1 

— 

2 

1 

1 

— 

3 

3 

3 

3 

1 

3 

3 

3 

3 

1 

=  Fe 

1 

1 

1 

1 

1 



3 

a 

3 

1 

1 

3 

3 

3 

3 



=  Fe 

— 



1 

1 

1 



3 

a 

2 

1 

1 

3 

3 

3 

3 

'  — 

=  Cr 

1 



1 

1 

1 



3 

a 

3 

1 

1 

3 

3 

3 

2 

— 

=  Al 

— 

— 

3 

1 

1 

3 

— 

a 

3 

3 

16 

3 

3 

— 

3 

1 

=  Sn 









.  

_ 



a 



. 



3 

3 







-Sn 









o 

3 





3 

3 



— 

3 

3 

— 

— 

=  Sb 

— 







2 









— 



— 

— 

— 

— 

.  — 

=  As 

1 

1 

3 

1 

1 

1 

3 

3 

3 

3 

1 

3 

3 

3 

3 

1 

=  Cd 

1 

1 

3 

1 

1 

3 

3 

3 

3 

3 

1 

3 

3 

3 

3 

1 

=  Cu 

2 



3 

1 

1 

3 

3 

O 

3 

3 

1 

3 

3 

3 

3 

1 

=  Bi 

1 

1 

3 

2 

4 

3 

3 

3 

3 

3 

1 

3 

3 

3 

3 

1 

=  Pb 

1 

— 

•  — 

— 

1 

2 

3 

— 

3 

2 

1 

3 

3 

3 

3 

1 

=  Hg 

1 

— 



.  _ 

3 

3 

.  



3 

3 

1 

3 

3 

3 

>  3 

1 

-Hg 

1 

2 

3 

2 

2 

3 

3 

3 

3 

3 

1 

3 

3 

3 

3 

1 

—  Ag 

—  =  Unknown,  rare,  or  not  readily  formed.  1  =  Keadily  soluble  in  water. 
2  =  Sparingly  soluble  in  water;  generally  dissolve  readily  in  very  dilute  acids.  3  = 
Practically  insoluble  in  water,  but  dissolve  readily  in  4N  mineral  acids.  4  =  Eequire 
hot,  strong  acids  for  appreciable  solvent  effect.  5  =  Practically  insoluble  in  acids. 

a  —  Hydroxid  is  precipitated.     6  —  Insoluble  basic  salt  likely  to  form. 


INDEX. 


NOTE; — /  indicates  frontispiece. 


Acetates,  122,  123 

action  of  heat,  122 

amyl,  303 

basic,  118,  122,  194,  226,  264,  265, 
303 

detection,  302,  303 

ethyl,  23,  24,  303 

solubility,  122 

stability,  122 
Acetic  acid,  122 

detection,  302,  303 

repression  of  ionization  of,  31,  264, 

268,  280,  281 
Acetylene,  120 
Acid  Analysis,  275-309 

comparison  with  Basic  Analysis,  242, 

275 
Acid  forming  elements,  15,  174 

properties  due  to  hydrogen  ions,  15 

properties    increased    by    oxidation, 

45,  175 
Acids,  action  on  bases,  19 

action  on  indicators,  15 

action  on  litmus,  15 

action  on  metals,  16 

action  on  oxids,  20 

anhydrid,  42,  54 

decomposition  on  heating,  54 

definition,  15 

dissociation,  16,  18 

furnish  hydrogen  ions,  15,  21,  174 

list  of,  16,  275 

meta,  175 

monobasic,  etc.,  17 

organic,  121 

ortho,  175 

polybasic,    17 

preliminary  tests,  276 

pyro,   175 


stability,   54 

strong  and  weak,  15,  16 

weak  displaced  by  strong,  35 
Active  agent,  meaning  of,  59 
Agate,  133 

Aids  to  chemical  reactions,  7,  8 
Alcohol,  amyl,  120,  270,  271,  302,  303 

ethyl  or  ordinary,  23,  24,  25,  74,  120, 

123,  235,  258,  259,  262,  263, 

270,   271,   284,   285,   302,   303 

as  reducing  agent,  47,  48,  262, 

263 

meaning  of  term,  120,  123 

methyl  or  wood,  114,  120,  282 
Alkali,  detection,  242 

fixed,  66 

group,    65,   66 

metals,  65,  66 

volatile,  66,  75 
Alkaline  earths,  90,  91 
Alloys,  analysis  of,  240 

Devarda's,  207,  300,  319 

mercury,  106 

platinum,  236 
Alum,  14,  117 

chrome,   193 

ferric,   225 

manganese,  214 
Alumina,  115 

Aluminates,  115,  116,  117,  266,  267,  301 
Aluminum,  115 

action  with  alkalies,  115,  301 

blowpipe  reaction,  115 

detection,  266,  267 

group,  111,  112 

hydrolysis  of  salts,  115,  118 

hydroxid,  116 

ions,  116 

properties  of  metal,  115 

relation  to  chromium  and  iron,  115 


334 


INDEX. 


salts,  115-113 
Amalgams,  106 
Amethyst,  133 
Ammonia,  146,  147 
complex  ions,  147 
oxidation,  39 

specific  gravity  of  solutions,  323 
water,  76,  147 
Ammonium,  75 

acetate  as  group  reagent,  264 
carbamate,  77,  270 
carbonate,  76,  77 

as  group  reagent,  264 
preparation   of   reagent,   318 
chlorid,  solvent  action,  113,  114,  125, 

200,  277,  279,  281 
detection,  76,  249 
hydroxid,  75,  76,   147 

repression  of  ionization,  31,  264, 

265,  268,  269 

specific  gravity  of  solutions,  323 
ion,  75 
molybdate,  78 

as  reagent  for 

arsenic  acid,  164,  287 
phosphoric    acid,    156,    260, 

261,  286,  287 

preparation  of  reagent,  318 
polysulfid,  77 

as  group  reagent,  252,  253 
as  reagent  for 

antimony,     167,     168,     169, 

252,  253 

arsenic,  162,  164,  252,  253 
cyanids,  125,  126,  294,  295 
gold,  90,  252,  253 
platinum,  235,  252,  257 
tin,  139,  141,  252,  257 
preparation    of    reagent,    318 
properties,  75,  147 
salts,  75-78 
sulfid,  77 

as  group  reagent,  268 
preparation  of  reagent,  318 
Amyl  acetate,  303 

alcohol,  120,  270,   271,  302,   303 
Analysis,  3 

acid,  275-309 
basic,  248-273 


chemical,  3 

gas,  243 

microchemical,   246 

of  alloys,  240 

of  insoluble  substances,  241 

of  organic  substances,  308,  309 

of  partial  solutions,  240 

qualitative,  3 

quantitative,  3 

systematic,  239 
Anhydrid,  definition,  42 
Anhydrous,  definition,  33 
Anion,   definition,   10 
Anions,  tests  for,  275-309 
Anode,  definition,  10 
Antimonates,  169 
Antimony,  165 

acids,  167,  168,  169 

blowpipe  reactions,  165,  166 

butter  of,  167 

comparison  with  arsenic,   166 

comparison  with  bismuth,  172 

detection,   165,  256,  257 

flame,  165 

ions,  166,  167 

pentachlorid,  168 

preparation  of  solution,   318 

pentavalent,  165,  168 

properties  of  element,  165,  257 

salts,  166 

stain  on  porcelain,  166 

trichlorid,  167 

preparation  of  solution,  318 

trivalent,  165,  166-168 
Appendix,  313-331 
Aqua  regia,  43 

as  solvent,  43,  89,  111,  234,  235,  240, 

250,  251,  252,  253,  258,  259 
Argon,  63,  65 
Arrow,  meaning  of,  5 

double-headed,  8 
Arsenates,  163,  164 

detection,  288,  289 

solubility,  163 
Arsenic,  156-159 

acid,  163 

detection,  288,  289 
salts,  164 

action  of  acids,  156 


LJSIDEX. 


335 


blowpipe  reaction,  157 

comparison  with  antimony,  166 

detection,  158,  256,  257 

flame,  158 

hydroxid,  159,  161 

ions,  157 

mirror,  158 

pentachlorid,  preparation  of  solution, 
318 

pentasulfid,   160 

oxidation  of,  50,  256,  257 

pentavalent,   157,  160 

salts,  159,  160 

stain  on  porcelain,  158 

trichlorid,  160 

preparation  of  reagent,  318 

trivalent,  157,  159,  160 

white,  156 
Arsenious  acid,  161 

detection,  288,  289 

salts,  161,   162 
Arsenites,  161,  162 

detection,  286,  287 

solubility,   161 

Arseno-molybdates,  164,  287 
Arsin,  158 
Atom,  definition,  8 

weight  of,  57,  326 
Atomic  weights,  326 
Aurates,  90 
Auric  salts,  89 
Aurous  salts,  89 


Barium,  98 

carbonate  as  reagent,  118,  225 

chlorid  as  reagent,  276,  278,  290 

detection,  270,  271 

dioxid,  178 

flame,  98,  270 

ion,  99 

properties  of  metal,  98 

salts,  98-101 

spectrum,  /,  98 

table  of  salts,  101 
Baryta  water,  99 
Base-forming  elements,  18,  174 
Bases,  definition,  18 

decomposition  on  heating,  54 


dissociation,  18 

furnish  hydroxyl  ions,  18,  21,  174 

monoacid,   etc.,  20 

polyacid,  20 

removal,  242 

strong  and  weak,  18 
Basic  Analysis,  248-273 

comparison  with  Acid  Analysis,  242 

group  separations,  248 
Basic  properties,  18 

decreased  by  oxidation,  45 
Bead  tests,  borax,  56,  247,  268 

meta-phosphate,    78,    192,    260,    261, 
280 

Na2C03  +  KN03,   219 
Benzene,  120 
Beryllium,  62,  90,  91 
Bettendorf  test  for  arsenic,  157 
Bismark  brown,  149 
Bismuth,  169-172 

acid,  171 

blowpipe  reaction,  170 

chlorid,  preparation  of  reagent,  318 

comparison  with  antimony,  172 

detection,  254,  255 

ions,  170 

nitrate,  preparation  of  reagent,  319 

properties  of  metal,  169 

salts,   170-172 

solubility  of  salts,  170 

subnitrate,  170 

sulfate,  preparation  of  reagent,  319 
Black-faced  type,  meaning  of,  3 
Blowpipe  flame,  53 

reactions,   55,   56 
Boiling-points  of  elements,  326 
Bonds,  significance  of,  12,  40 
Bone-black,  119 
Borates,  113,  114 

action  of  heat,  55,  113 

detection,  282,  283 

methyl,  114 

precipitation,  113 

solubility,  113 

stability,    55,    113 
Borax,  56,  113 

bead  reactions,  56,  247 

glass,  113 
Boric  acid,  112-114 


336 


INDEX. 


anhydrid,  113,  114 

detection,  114,  282,  283 

salts,  113,  114 
Boron,  112 

flame,  112,  282,  283 

group,  111 

properties  of  element,  112 

salts,  112 

Bromates,  208,  209 
Bromids,  209,  210 

action  of  heat,  55,  209 

detection,  298,  299 

reaction  with  cone,  sulfuric  acid,  209 

solubility,  209 
Bromin,  208 

acids,  209 

as  oxidizing  agent,  41,  43,  48,  189, 
208,  290,  291,  297 

chlorid,  209 

detection,  209,  308,  309 

ions,  210 

properties,  208 

water,  208 
Brucin,  151,  206,  300,  301 

preparation  of  reagent,  319 
Bunsen  burner,  improvised,  114 

flame,  52 


Cadmium,  104 

blowpipe  reaction,  105 

detection,  254 

ion,  105 

properties  of  metal,  104 

salts,   104-106 
Caesium,  64,  65,  66 
Calcium,  94 

detection,  270,  271 

flame,  94,  270 

hydroxid,  95 

ion,  94 

phosphate,  35,  97 

properties  of  metal,  94 

salts,  94-97 

spectrum,  /,  94 
Caliche,  70 
Calomel,  111 
Carbids,  120 
Carbon,  119 


acids,  121 

amorphous,  119 

compounds,  120 

detection,  308,  309 

dioxid,  120 

disulfid,  120,  179,  296,  297,  298 

group,  119 

hydrids,  120 

monoxid,  119 

properties  of  element,  119 

removal,  241 
Carbonaceous  matter,  detection,  260,  261 

interferes  with  analysis,  260 

removal,  262,  263 
Carbonates,  122 

acid,  92,  95,  122,  222 

action  of  heat,  121 

basic,  69 

detection,  282,  283 

precipitation,  68,  242,  261,  271 

solubility,  66,  122,  271 

stability,  54,  121 
Carbonic  acid,  121 

anhydrid,  120,  122 

detection,  282,  283 

salts,  121,  122 

Carbonization  (see  Charring) 
Cerium,  119 
Chalcedony,  133 
Chalybeate  springs,  222 
Charcoal,  119 

reactions,  55,  246,  247 
Charring,   55,  124,  260,  261,  284 
Chemical  formulae,  6 

reactions,  aids  to,  7 
Chemistry,  definition,  3 

general,  4 
Chlorates,  206 

action  of  heat,  206 

detection,  304,  305 

formation,  201,  204 

reaction  with  cone,  sulfuric  acid,  206 

solubility,  206 
Chloric  acid,  205,  206 

as  oxidizing  agent,  42,  43,  205,  206, 
305,  306 

detection,  304,  305 

ion,  206 

reactions  with  brucin,  151,  206 


INDEX. 


reduction,  305,  306 

salts,  206 

"Chlorid  of  lime/'  204 
Chlorids,  203 

action  of  heat,  55,  203 

decomposition,  298 

detection,  298,  299 

reaction  with  cone,  sulfuric  acid,  203 

separation,  298 

so,  ibility,  203,  251 
Chlorin,  200 

acids,  202,  203,  205,  207 

as  oxidizing  agent,  41,  43,  48,  201, 
296,  297,  298,  299 

detection,  202,  308,  309 

dioxid,  205 

distinction   from   hypochlorous   acid, 
202 

group,  198 

heptoxid,  207 

ions,  200,  201 

monoxid,  203 

oxids,  201,  203,  205,  207,  304 

properties,  200,  201 

reactions  of  ion,  13,  14 

water,  200,  204 
Chlorites,  205 
Chloro-chromic  anhydrid,  196,  197,  298, 

299 

Chloroform,  298,  299 
Chlorous  acid,  205 
Chlorplatinates,  235 
Chromates,  194-197 

action  of  heat,  192 

detection,  280,  281 

precipitation,  195 

solubility,  196 
Chrome  alum,  193 

iron  ore,  190 
Chromic  acid,  194,  195 

anhydrid,  194,  195 

detection,  280,  281 

ion,  194 

relation  to  dichromic  acid,  195 

salts,  196,  197 
Chromites,  193 
Chromium,  190 

acids,  192,  193,  194-197 

blowpipe  reaction,  192 

23 


337 


borax-bead  test,  192 

compounds,  191 

detection,  192,  266,  267 

divalent,  191 

group,  190 

heptavalent,  192 

hexavalent,  192,  194-197 

hydroxids,  46,  193 

ions,  192,  194,  197 

properties  of  element,  190 

salts,  192-194 

tests,  192 

trivalent,  191,  192-194 
Classification  of  elements,  62 
Closed-tube,   53 

reactions,  53-55,  107,  157,  243,  244 
Cobalt,  227 

acetate,    reagent    for    cyanates    and 
thiocyanates,  128,  305,  306 

blowpipe  reaction,  228 

borax-bead  reaction,  228 

compounds,  227-231 

detection,  268,  269 

ions,  228,  229 

nitrate  reactions,  245 

with   aluminum    oxid,   115,   245, 

266 

with  silicates,  245 
with  zinc  oxid,  102,  245 

properties  of  metal,  227 

salts,  228 

separation  from  nickel,  231,  233,  234, 

268,  269 

Cobaltic  compounds,  228 
Cobaltinitrite,  231 
Cobalticyanids,  230,  233 
Cobaltocyanids,  230 
Cobaltous  compounds,  228,  229-231 
Coke,   119 

Colloidal  precipitates,  37,  159,  253,  269 
Combustion,  173 
Comparative  solubilities,  34 

strengths  of  acids,  16 

strengths  of  bases,  18 
Complex  ions,  13,  14 

decomposition  of,  14 

organic  ions  of  aluminum,  117,  260 
chromium,  193,  260 
copper,  81 


338 


INDEX. 


iron,  224,  260 
Concentration,  definition,  23 

how  expressed,  25 
Conversion  factors,  329 
Copper,  79 

action  of  acids,  79 

as  reagent  for  As,  159 

blowpipe  reaction,  80 

detection,  254,  255,  256 

flame,  80 

group,  79 

ions,  80 

precipitation,  79,  80 

properties  of  metal,  79,  80 

salts,  80 
Copperas,  223 
Corrosive  sublimate,  109 
Crystal,  definition,  33 
Crystallization,  definition,  33 
Cupric  compounds,  81-83 

detection,  254 

ion,  80,  81 

salts,  80,  81-83 
Cuprous  ion,  80 

salts,  80,  81,  83,  84 
Cyanates,  128 

detection,  304,  305 
Cyanids,  126 

action  of  heat,  55,  126 

as  reagent  for  As,  157 

as  reducing  agents,  126 

complex,  127 

decomposition,  126,  242 

detection,  242,  294,  295 

hydrolysis,  126 

insoluble,  126,  127 

of  metals  with  valence  greater  than 
two,  126 

solubility  of,  126 
Cyanogen,  125 


Decantation,  washing  by,  38 
Decomposition,  of  complex  ions,  13,  14 

on  heating,  54-56 

with  concentrated  sulfuric  acid,  55 
Deliquescence,  33 
Desiccation,  134,  262,  263 
Determining  factors,  26,  27 


Devarda's  alloy,  207,  300,  319 
Diamond,  119 
Dichromates,  195,  196 

detection,  280,  281 

ion,  194 

precipitation  with,  195 

reaction  with  cone,  sulfuric  acid,  196 

solubility,  196 
Dichromic  acid,  195 

anhydrid,  195 

as  oxidizing  agent,  50,  253,  262,  263, 
296,  297,  298,  299 

detection,  28JO,  281 

ion,  194 

relation  to  chromic  acid,  195 
Dissociation,  8 

affected  by  common  ions,  31 

by  steps,  18 

degree  of,  10,  25 

electrolytic,  10 

equilibrium,  25 

in  solution,  9 

of  acids,  16,  18 

of  bases,  18 

of  salts,  20 

of  water,  29 

on  heating,  8,  56 

solvents  causing,  13 
Distillate,  definition,  36 
Distillation,  definition,  36 
"Dry  way"  reactions,  5 
oxidations,  52 


Efflorescence,  33 

Electricity,  aids  chemical  reactions,  8 

conduction  in  solution,  10 

on  ions,   9,   11 

unit  charge  on  ions,  11 
Electrolyte,  definition,  10 
Electrolytic    dissociation,    10     (see    also 

"lonization") 
Electrolysis,  256,  257 

definition,  10 

explanation,    10,    257 
Elements,  4 

acid-forming,   15 

base-forming,  18 

classification,  62 


INDEX. 


339 


list  of,  324,  326 

metallic,  18,  174 

non-metallic,  15,  174 

properties  changed  by  oxidation,  45 

symbols  for,  5 
Epsom  salt,  93 
Equation,  definition,  5 

should  be  balanced,  6 
Equilibrium,  22-28 

changes  with  temperature,  27 

definition,  23 

equation,  24 

law,  23 

reaction,  23 
Equivalents,  definition,  58 

system  of,  58 

values  of,  59 
Erbium,   65 

Ester,  meaning  of  the  term,  121 
Ether,  284 

meaning  of  term,  120 
Ethyl  acetate,  23,  24,  303 

alcohol,  23,  24,  25,  74,  120,  123,  235, 
258,  259,  262,  263,  270,  271,  284, 
285,  302,  303 

Evaporation,  definition,  36 
Examination,  preliminary,  243-247 

F 

Factors,  conversion,  329 
determining,  26,  27 
Ferrates,  221 

Ferric,  chlorid,  hydrolysis  of,  30,  225 
compounds,  221,  224-227 

as  oxidizing  agents,  221,  225 
detection,  227,  249 
hydrolysis  of,  221,  264,  265,  303 
ion,  224-227 

reagent    for    ferrocyanids,    131, 

226,  249,  264 
reagent    for    thiocyanates,    129, 

227 

salts,  221,  224-227 
Ferricyanic  acid,  131,  226 
detection,  292,  293 
salts,  131,  132 
Ferricyanids,  131,  132,  226 
detection,  292,  293 


reagent   for   ferrous   ions,   131,   226, 

249 

Ferrocyanic  acid,  130,  223 
detection,  292,  293 
salts,  130,  131 

Ferrocyanids,  130,  131,  223 
action  of  heat,  130 
detection,  292,  293 
formation,    127 
reagent  for  ferric  ion,  131,  226,  249, 

264,  295,  308,  309 
solubility,   130 
Ferrous  compounds,  221,  222-224 

as  reducing  agents,  47,  221,  258, 

259 

detection,  227,   249 
oxidation,  292 
with    NO,    147,    149,    151,    300, 

301,  302,  303 
ion,  222-224 

reagent    for    ferricyanids,    131, 

22'6,  249 

salts,  221,  222-224 
Film  tests,  246 

on  charcoal,  246 
on  porcelain,  158,  166,  246 
Filter,  definition,  37 

fitted,  38 
Filtrate,  definition,  37 

cloudy,  how  avoided,  37 
Filtration,  37 
"Fireworks,"  73,  206 
Flame,  blowpipe,  53 
Bunsen,  52 
coloration,  56,  247 
oxidizing,  53 
reducing,  53 
Flash-light  powder,  91 
Flint,  133 
Fluorids,  198-200 

action  of  heat,  55,  199 
detection,  278 

reaction  with  cone,  sulfuric  acid,  199 
solubility,  199,  200 
Fluorin,  198 
acid,  199 

action  on  water,  177,  198 
detection,  278 


340 


INDEX. 


group,  198 
ion,  136,  198,  200 
properties,  198 
Fluospar,  97,  200 
Fluosilicic  acid,  136 
Formulae,  chemical,  6 
Fusion,  with  ammonium  nitrate,  78,  122, 

124,  262 

with   calcium  carbonate   and   ammo- 
nium  chlorid,  242 

with  potassium  acid  sulfate,  188,  242 
with     sodium     and     potassium     car- 
bonates, 69,  242,  262 
with  sodium  carbonate  and  potassium 

nitrate,  192,  216,  242 
with  sodium  hydroxid,  241,  294,  295 

Q 

Gadolinium,  65 
Gallium,  118 
Gas  analysis,  243 

carbon,  119 

Gaseous  substances,  as  determining  fac- 
tors, 26 

how   indicated,   3 

liberation  of,  240,  244,  245 

prediction  of  formation,  26 
Germanium,  137 
Glass,  etching  of,  199,  278,  279 

composition,   199 
Glauber's  salt,  70 
Gold,  89 

blowpipe  reaction,  89 

detection,   258,   259 

fulminating,  90 

mosaic,  241 

properties  of  element,  89 

salts,  89 

Gram-molecular  solutions,  58 
Gram-molecule,  definition,  57 
Graphite,  119 
Gypsum,  96 


Halogen  group,  198 

detection,  308,  309 
Hardness  of  water,  96 
Heat,  action  on  acids,  54 


on  bases,  54 

on  salts,  54,  55,  56 

aids  chemical  activity,  7 

causes  dissociation,  8,  56 

of  neutralization,  19 
Heating  in  borax-bead,  56,  247 

in  closed-tube,  53-55,  107,  157,  243, 
244 

on  charcoal,  55,  246,  247 

with  concentrated  sulfuric  acid,  55, 

245 

Helium,  62,  65 

High-temperature  reactions,  54-56,  150 
Hydration,  of  acids,  175 

of  antimonates,  168,  169 

of  antimonites,  167 

of  arsenates,  163 

of  arsenites,  161 

of  borates,  113 

of  chromates,  194,  195 

of  phosphates,  153,  154 

of  silicates,  133,  134 

of  sulfates,  186,  187 
Hydriodic  acid,  211 

as  reducing  agent,  48,  49,  211,  295, 
296 

detection,  296,  297 

ions,  212 

salts,  212 
Hydrobromic  acid,  209 

as  reducing  agent,  48,  209,  298 

detection,  298,  299 

ions,  210 

salts,  209,  210 
Hydrochloric  acid,  202 

as  group  reagent,  250 

as  reducing  agent,  48,  49,  202 

as  solvent,  239 

detection,  298,  299 

ion,  203 

salts,  203 

specific  gravity  of  solutions,  322 
Hydrocyanic  acid,  125,  126 

detection,  128,  294,  295 

salts,  126-128 

separation,  295 
Hydrofluoric  acid,  198,  199 

action  on  glass,  199 


INDEX. 


341 


detection,  178,  279 
ion,  136,  200 
salts,  199,  200 
Hydrogen,  bromid,  209 
chlorid,  202 
cyanid,    125,    126 
detection,  308,  309 
dioxid,  117 

as  oxidizing  agent,  42,  43,  177, 

178 

as  reagent  for  chromium,  192 
detection,  306,  307 
salts,  178 
iodid,  211 

ion,  properties  of,  15 
nascent,   40 

as  reagent  for  antimony,  166 

as  reagent  for  arsenic,  158 

as  reducing  agent,  158,  166,  256, 

257,  300,  301 
oxid   (see  Water),  176 
peroxid  (see  dioxid),  177 
replaced  by  metal,  17 
sulfid   (see  also  Hydrosulfuric  acid), 

180 

action  on  metals,  184 
as  group  reagent,  252 
as  reducing  agent,  47,  48,  50, 

180,  225,  252,  253,  265 
detection,    183,    274,    275,    290, 

291,  294,  295 
oxidation,  180,  253,  265 
solubility,  180 
water,  180,  188 
Hydrolysis,  30,  264,  265 
Hydronitric  acid,  146 
Hydrosilicofluoric  acid,  136 
detection,   278,    279 
salts,  136,  137 
Hydrosulfuric  acid    (see  also  Hydrogen 

sulfid),    180 
ions,  180 
salts,  180-184 
Hydroxids,  acting  either  as  acid  or  base, 

22,  175 

aluminum,  116 
antimony,  166 
arsenic,  159,  161 
chromium,  193 


gold,  90 
lead,  143,  144 
platinum,   235 
stannous,  139 
stannic,  140 

insoluble,  35 

of  metallic  elements,  18,  68,  174 

of  non-metallic  elements,  15,  174 

organic,  120 

precipitation  of,  68 
Hygroscopic,  meaning  of  term,  33 
"Hypo,"   190 
Hypobromites,  208,  209 
Hypochlorites,  203-205 

action  of  heat,  204 

detection,  304,  305 

formation,  201 

solubility,  204 
Hypochlorous  acid,  203,  204 

action  on  hydrogen  chlorid,  204 

anhydrid,  203 

as  oxidizing  agent,  42,  43,  204 

detection,  304,  305 

ion,  205 

salts,  204,  205 
Hypoiodites,  211 
Hyponitrous  acid,  147 

anhydrid,  147 
Hyposulfurous  acid,  185 


Indium,  118 

spectrum,  / 

Infusible  substances,  245 
Inorganic  substances,  definition,  13 

dissociate  generally,  13 
Insoluble,  hydroxids,  35 
salts  of  weak  acids,  35 
substances,  analysis  of,  241 
as  determining  factor,  27 
meaning  of  term,  34 
prediction  of  formation.  26 
solution  of,  36,  241,  242 
Interfering  substances,  260-263 
detection,  260,  261 
removal,  262,  263 
lodates,  211,  299 
lodids,  212 

action  of  heat,  55,  212 


342 


INDEX. 


detection,  296,  297 

reaction  with  cone,  sulfuric  acid,  212 

removal,  283 

solubility,  212 
lodin,  210 

acids,  211 

as  oxidizing  agent,  41,  43,  48,  189, 
210,  283,  290,  291 

chlorid,  211,  297 

detection,  211,  308,  309 

ions,  210 

properties,  210 

removal,  283 

stains,  210 

water,  210 

preparation,  319 
Ions,  addition  of  common,  31 

complex,  12,  13 

decomposition  of,  13,  14 

definition,  9 

electricity  on,  9,   11 

how  indicated,  11 

metallic,  10 

non-metallic,  10 

potential,  28 

reactions  of,  13 

valence,  11 
lonization,  definition,  9 

and  equilibrium,  25 

how  affected  by  common  ions,  31 

of  acids,  16,  18 

of  bases,  18 

of  salts,  20 

repression  of,  31 
Iridium,   220 
Iron,  220 

acids,  221 

blowpipe  reaction,  221 

borax-bead  reaction,  221 

detection,  227,  249 

group,  220 

ions,  221 

tests  for,  227 

oxids,  221 

properties,  220 

rust,  220,  222 

salts,  221-227 
Italics,  meaning  of,  3 


Kathion,   definition,   10 
Kathions,  tests  for,  248-273 
Kathode,  definition,  10 
Krypton,  65 


Lampblack,  119 
Lanthanum,  111 
Lead,  142 

acetate  as  reagent  for  hydrogen  sul- 
fid,  252,  294,  308,  309 

blowpipe    reaction,    142 

detection,  250,  251,  254 

group,  137 

hydroxid,  143 

ions,  143,  144 

properties  of  metal,  142 

red,  143 

salts,  143-145 

sugar  of,  143 

white,   144 

Light  aids  chemical  reactions,  7 
Lime,    95 

salts,  95-97 

slacked,  95 

water,  95,  282,  283,  308 
Liquids,  analysis  of,  243 

how  indicated,  3 
Litharge,  142 
Lithium,  66 

detection,  272,  273 

flame,  66,  272,  273 

ion,   66 

properties  of  metal,  66 

salts,  66 

spectrum,  /,  66,  272 

M 
Magnesia,  91 

mixture,  92 

preparation,  319 
Magnesium,  91 

detection,  272,  273 

group,  90,  91 

ion,  92 

properties  of  metal,  91 

salts,  92-94 

solubility  of  salts,  92,  289 


INDEX, 


343 


Manganese,  213 

acids,  214,  215,  218 
alum,  214 

borax-bead  reaction,  216,  268 
chlorid,  preparation  of  reagent,  319 
detection,  268,  269 
dioxid,  178,  214 
divalent,  213,  216-218 
group,   212 

heptavaleut,   216,  218-220,  268,  269 
hexavalent,  215,  268,  269 
hydroxids,  45,  216 
ions,  216 

oxidation,  216,  268,  269 
oxids,  213,  214,  216 
properties  of  element,  213 
salts,  213,  214,  216-218 
tetravalent,  214,  277 
trivalent,  214,  277 
Manganates,  215,  268,  269 
Manganicyanids,  218 
Manganites,  214,  216,  217 
Manganocyanids,  218 
Marsh  test,  158.  166 
Mass  action,  definition,  25 

discussion  of,  23-27 
illustration  of,  28-30 
law  of,  23,  24 
Matches,  73,  206 
Match  test,  246 
Measures,  English,  328 

metric,   328 

Mechanical  force  aids  reactions,  7 
Melting-points  of  elements,  326 
Mendeleeff,  63 
Mercuric,    chlorid    as    reagent    for   stan- 

nous  salts,  138,  256,  257 
compounds,  107 

action  of  NH3  on,  108 
detection,  254 
reduction,  107,  110 
ion,  107 
salts,  107-110 
Mercurous  compounds,  107 

action  of  NH3  on,  110,  111 
detection,  250 
reduction,  107,  111 
ions,  107 
nitrate,  preparation  of  reagent,  319 


as  reagent  for  ammonia,  304,  305 
as  reagent  for  cyanids,  294,  295 

salts,   110 
Mercury,  106 

detection,  107,  250 

halids,  dissociation  of,  21 

in  separation  of  chlorin  from  hypo- 
chlorites,  202 

ions,  107 

properties  of  metal,  106 

salts,  107 
Meta-acids,  175 
Metal,  added  to  acid,  16 

definition,  174 
Metals,  analysis  of,  240 
Metathesis,  3 

Methyl  alcohol,  lj.4,  120,  282 
Microcosmic  salt,  78 
Minus  ( — )  sign,  meaning  of,  11 
Mixtures,  4 
Molecules,  definition,  8 

weight   of,   57 
Molybdenum,  65,  190 

N 

Naphthalene,  120 
Naphthylamin,  149 
Nascent  condition,  40 

hydrogen    (see  Hydrogen) 
Neodymium,  64 
Neon,  62,  65 
Nessler  reagent,  109 
Neutralization,  19,  20 

heat  of,  19 
Nickel,  231 

blowpipe  reaction,  232 

borax-bead  reaction,   232 

compounds,  231-234 

detection,  268,  269 

ions,  232 

properties  of  metal,  231 

salts,  232 

separation    from    cobalt,    231,    233, 

234,   268,   269 

Nickelic  compounds,  232,  233,  268,  269 
Nickelous  compounds,  232-234 
Niobium,  172 
Nitrates,  150 

action   of  heat,  55,  150 


344 


INDEX. 


decomposition,  55,  150 

detection,  300,  301 

solubility,  150 
Nitric  acid,  149-151 

anhydrid,  148 

as  oxidizing  agent,  39,  43,  50,  150, 
151,  239,  240,  255,  261,  264,  265, 
268,  269,  286,  287,  300,  301,  308, 
309 

as  solvent,  239 

detection,  300,  301 

fuming,  150 

ion,  151 

reduction,  151,  300,  301 

salts,  150 

specific  gravity  of  solutions,  323 

tests  for,  151,  300,  301 
Nitrids,  146 
Nitrites,  148,  149 

action  of  heat,  148 

detection,  149,  302,  303 

solubility,   149 
Nitrogen,  146 

compounds,  146-151 

detection,  146,  308,  309 

dioxid,  147,  149,  151 

group,  145 

ions,  146 

monoxid,  147 

oxids,  147-148,  173 

pentavalent,  146,  149 

pentoxid,  148 

tetroxid,  148,  150 

trioxid,  148 

trivalent,  146,  148 
Nitrous  acid,  148,  149 

anhydrid,  148 

as  oxidizing  agent,  149,  296,  297,  302 

decomposition,  148,  301 

detection,  149,  301,  302,  303 

either  oxidizing  or  reducing  agent, 
149 

ions,  149,  231 

solubility,  297 
Nomenclature,  324-325 
Non-dissociation,  as  determining  factor, 

27 
Non-metallic  elements.  174 


Normal  salts,  22 
solution,  59 


Onyx,  133 

Organic  acids,  121 
chemistry,    120 
hydroxids,  120 
oxids,  120 
salts,  121 
substances,  13,  120-132 

action  of  heat,  122,  124,  261 
analysis  of,  308,  309 
definition,  13 
detection,   260,  261 
do  not  dissociate,  13 
interfere  with  analysis,  260 
reduce   hot    cone,    sulfuric    acid, 

124 

removal,  260,  261 
Ortho-acids,  175 
Osmium,  220 
Oxalates,  123,  124 

action  of  heat,  123 
detection,  284,  285 
solubility,  123 
Oxalic  acid,  123 

as  reducing  agent,  47,  48,  49,  123 
decomposition,  123 
detection,  260,   261,  284,  285 
interferes  with  analysis,  260,  261 
removal,  262,  263 
salts,  123,  124 
Oxidation,  definition,  38 

increases  acid  properties,  45,  175 
increases  charge  on  kathions,  52 
in  "dry  way,''  52 
Oxidizing  agents,  acids  as,  42,  43 
definition,   39 
direct  addition   of,  41 
elementary  substances  as,  41 
halogens  as,  41 
in  analysis,  239,  240,  252 
list  of,  41 

oxygen  equivalent  of,  43 
reactions  with  reducing  agents,  49- 

51 

removal  of,  239,  240,  283 
salts  as,  42 


INDEX. 


345 


tests  for,  276,  277 

value  of,  43,  49 
Oxidizing  flame,  53 
Oxids,  174 

form  hydroxids,  174 

how  formed,  174 

ignited,  174,  242 

insoluble,  35,  174,  242 

organic,  120 

volatility  of,  54 
Oxygen,  compounds,  174-178 

detection,  173 

group,  172,  173 

ion,  176 

properties,  173 

test  for,  173 
Ozone,  174 


Palladium,  220 

Perchlorates,  207,  208 

action  of  heat,  207,  306,  307 
detection,  306,  307 
formation,  206 
solubility,    208 

Perchloric  acid,  207 
anhydrid,   207 
detection,  306,  307 
ion,  208 
salts,  207,  208 

Perchromic  acid,  192,  266 

Periodates,  211 

Periodic  law,  63 

Periodic  System,  63-65 
Group  0,  65 
Group  1,  first  type,  65 
Group  1,  second  type,  79 
Group  2,  first  type,  90 
Group  2,  second  type,  100 
Group  3,  first  type,  111 
Group  3,  second  type,  118 
Group  4,  first  type,  119 
Group  4,  second  type,  137 
Group  5,  first  type,  145 
Group  5,  second  type,  172 
Group  6,  first  type,  172 
Group  6,  second  type,  190 
Group  7,  first  type,  198 
Group  7,  second  type,  212 


Group  8,  220 
Periodids,  210 
Permanganates,  219,  220 

action  of  heat,  219 

detection,  306,  307 

formation,  215 

solubility,  219 
Permanganic  acid,  218,  268,  269 

anhydrid,  216 

as  oxidizing  agent,  49,  123,  124,  219, 
220,  276,  277,  284,  298,  299,  306 

detection,  306,  307 

ion,  220 

salts,  219,  220 
Persulfuric  acid,  190 
Phenylenediamin,  149,  302 

preparation  of  reagent,  319 
Phosphates,  153-156 

action  of  heat,  55 

detection,  260,  261,  286,  287 

of  alkali  metals,  70 

precipitation,  155 

removal,   156,  262,   263 

solubility,  66,  97,  154,  155 

stability,  55 
Phosphin,  152 

Phospho-molyMates,  156,  286,  287,  308 
Phosphonium  compounds,  152 
Phosphoric  acid,  153,  154 

anhydrid,  153 

detection,  260,  261,  286,  287 

glacial,  154 

interferes  with  analysis,  260,  261 

ions,  155 

meta,  154 

ortho,  153,  154 

pyro,  154 

removal,  156,  262,  263 
Phosphorous  acid,  153 
Phosphorus,  151,  152 

black,  151 

compounds,   152 

detection,  152-156,  308,  309 

hydroxids,  152 

in  alloys,  240 

ions,  153 

oxids,  153 

pentavalent,  152 

red,  152 


346 


INDEX. 


trivalent,  152 

white,  151 

yellow,  151,  152 
Physical  properties,  4 

of  elements,  326 
Plaster  of  Paris,  96 
Platinic-chlorids,  235,  258,  259 
Platinum,  234 

acid,  235 

black,  234 

blowpipe  reaction,  235 

detection,  258,  259 

divalent,  234 

group,  220 

ions,  235 

properties  of  element,  234 

reactions   at   high   temperature,   236 

salts,  234 

sponge,  234 

tetravalent,  234,  235 

ware,  care  of,  235 

wire,  cleaning,  56 
Plumbic  acid,  143 
Plumbites,  144 
Plumbous  acid,  144 
Plus  (  +  )  sign,  meaning  of,  11 
Polysilicic  acids,  133,  134 

-sulfids,   184 

ammonium,  77 

-thionic  acids,  189 
Potash,  72 

caustic,  72 

lye,  72 
Potassioscope,  57,  272,  273 

fluid,  319 
Potassium,  72 

detection,  272,  273 

in  presence  of  sodium,  57,  272, 
273 

flame,  72,  272,  273 

hydroxid,  72 

ion,  72 

properties  of  metal,  72 

salts,  72-74 

table   of,   74 

spectrum,  /,  72,  272 
Potential  ions,  28 

series,  327 
Praseodymium,  64 


Precipitate,  character  of,  33 

colloidal,  37,  160,  253,  269 

definition,  37 

washing  of,  38 
Preliminary  examination,  53-57,  243-247 

tests  for  acids,  276 
Prussian  blue,  131,  226 
Prussic  acid,  125 

Pseudo-solution,  37,  159,  224,  253,  269 
Pure  substances,  4 
Purple  of  Cassius,  89,  259 
Pyro-acids,  175 


Qualitative  analysis,  3 

depends  on  comparative  solubili- 
ties, 34 

grouping    of    elements   in,    34 
method  of,  5,  34 
Quartz,  133 

B 

Radicles,  definition,  5 
Radium,  90,  91 
Reactions,  aids  to,  7,  8 

between  ions,  13 

blowpipe,  55,  56 

closed-tube,  53-55,  243,  244 

definition,  5 

"dry  way,"  5 

high-temperature,  54-56 

how  affected  by  temperature,  7 

in   borax-bead,   56,   247 

ionic,  50-52 

of  qualitative  analysis,  28 

on   charcoal,  55,  246,  247,  278,  290 

predicted,  26 

reversible,  8,  23 

"wet  way,"  5 

with  cobalt  nitrate,  102,  115,  245 

with  cone,  sulfuric  acid,  55,  123,  124, 
126,  130,  131,  183,  196,  199,  203, 
206,  209,  212,  245,  260,  278,  279, 
280,  298,  299,  300,  302,  303,  304, 
305 
Reagents,  definition,  5 

laboratory,  60,  313-320 

solid,  list  of,  321 
Reducing  agents,  decomposition  of,  47 

definition,  39 


INDEX. 


347 


list  of,  46 

oxygen  absorbed  by,  47 

reactions  with  oxidizing  agents,  49- 
51 

test  for,  276,  277 

value  of,  49 
Reducing  flame,  53 

Reduction,  decreases  charge  on  kathions, 
52 

definition,  39 

of  elements  from  compounds,  80 

reduces  acid  properties,  46 
Reinsch  test  for  arsenic,  159 
Repression  of  ionization,  31 

of  acetic  acid,  31,  264,  268,  280,  281 

of  ammonium  hydroxid,  31,  264,  265, 

268,  269 

Residue,   definition,  37 
Reversible  reactions,  23 
Rhodium,  220 
Rinnmann's  green,  102 
Rochelle  salt,  73 
Rubidium,  65,  66 
Ruthenium,  220 

S 

Salt,  common,  70,  198 

Epsom,   93 

Glauber's,   70 

microcosmic,  78 

Eoehelle,  73 

Sehlippe's,  169 

sea,  70,  198 

tin,  140 
Saltpeter,  72 

Chili,  70 
Salts,  acid,  21 

balanced,  28 

basic,  21 

decomposition  on  heating,  54 

definition,  18,  21 

dissociation  of,  20 

double,  14 

formed  by   strong  bases  with  weak 
acids,  29 

formed   by  weak  bases  with   strong 
acids,  30 

in  oxidation,  42,  44 

insoluble,  of  weak  acids,  35 


ionization  of,  20 

methods  of  formation,  20 

more  stable  than  acids,  42,  277 

most  stable  formed  by  most   stable 
acids,  54 

most     stable    formed    by    strongest 
bases,  54 

normal,  22 

organic,  121 

reactions  of  hydrolyzed,  30 

solubility  of,  22 

unbalanced,  28 
Samarium,  65 
Scandium,  111 
Schlippe's  salt,  169 
Sea  salt,  70,  198 

water  in  boilers,  93 
Selenium,  64,  172,  173 
Separation  of  solids  from  liquids,  36,  37 
Silica,  133 

jelly,  135 
Silicates,  134-136 

action  of  heat,  55,  134 

decomposition,  135,  136,  242 

detection,  280,  281 

hydrolysis,  135,  289 

solubility,  134 

stability,  55 
Silicic  acid,  133,  134 

anhydrid,  133,  134,  135 

detection,  260,   261,  280,  281 

interferes  with  analysis,  260 

ions,  134 

poly-,   134 

salts,  134 

separation,  262,  263 
Silicids,  132 
Silicofluoric  acid,  136 
Silicofluorids,  136,  137 
Silicon,  132 

compounds,  132,  279 

dioxid,  133 

fluorid,  132 
Silver,  84 

action  of  acids,  84 

blowpipe  reactions,  84 

chlorid,  86 

solubility  of ,  35,  85-87,  190,276 

detection,  250,  251 


348 


INDEX. 


halids,  86-88,  241 

decomposition,  298 
hydroxid,  85 
ions,  85 

reactions  of,  13,  14,  85-88 
mirror,  86,  125,  284,  285 
nitrate    as    reagent    for    acids,    276, 

277 
reagent  for  arsenates,  256,  257, 

258,  259,  288,  289 
reagent  for  chlorids,  298 
properties  of  metal,  84 
salts,  85-88 

solubility  of,  35,  85-87,  190,  276 
table  of,  88 
Slime,  how  avoided,  38 
Soda,  69 

caustic,  68 
lye,  68 
sal,  69 
washing,  69 
Sodium,  67 

action  on  water,  176,  177 
carbonate,  68,  69 

fusions  with,  69,  242 
hydrolysis  of,  29 
precipitation  by,  69,  242 
transposition  by,  68,  69,  242 
detection,  272,  273 
flame,  67,  272,  273 
hydroxid,  68 
ion,  67 

nitroferricyanid,  184,  308,  309 
nitroprussid,  184,  308,  309 
peroxid,  67,  178,  266,  267,  280,  298, 

299 

properties  of  metal,  67 
salts,  67-71 

table  of,  71 
spectrum,  /,  67,  272 
Solid  reagents,  list  of,  321 
Solids,  how  indicated,  3 
Solute,   definition,  32 
Solubilities,  comparative,  34 

table  of,  330,  331 

Solubility,   affected  by   temperature,   32 
definition,  32 
in  strong  acids,  35 
in  strong  bases,  36 


Solution  aids  chemical  reactions,  8 

conduction  of  electricity  in,  10 

gram-molecular,   58 

normal,  59 

of  ignited  oxids,  242 

of  insoluble  substances,  36,  241,  242 

of  silicates,  242 

original,  240 

phenomena  in  dilute —  14 

in  concentrated —  14,  28 

prepared,  242 

pseudo,  37 

saturated,  32 

saturated     with     precipitated     sub- 
stance, 33 

specific  gravity  of —  322,  323 

supersaturated,  32 

unsaturated,  32 
Solvent,  aqua  regia  as,  240 

definition,  32 

hydrochloric  acid  as,  239 

nitric  acid  as,  239 

water  as,  239 

Solvents,  causing  dissociation,  13 
Soot,  119 
Specific  gravity  of  elements,  326 

of  solutions,  322,  323 
Spectra,  table  of,  / 
Spectrum  of  a  barium,  /,  98 

caesium,  / 

calcium,  /,  94 

indium,  / 

lithium,  /,  66,  272 

potassium,  /,  72,  272 

rubidium,  / 

sodium,  /,  67,  272 

strontium,  /,  97 

thallium,  /, 
Spelter,  101 
Stannates,  139,  140 
Stannic,  acid,  137,  140,  141 

chlorid,  preparation  of  reagent,    J19 

compounds,  138 
detection,  249 
reduction,  256,  257 

hydroxid,  140,  141 

ions,  140-142 

salts,  140-142 
Stannites,  139 


INDEX. 


349 


as  regard  for  bismuth,  171,  254 
Stannous,  acid,  139 

chlorid  as  reagent  for  arsenic,  157 
as  reagent  for  mercury,  110,  111 
preparation  of  reagent,  319 

compounds,  138,  139,  140 

as  reducing  agents,  47,  251,  259, 

283,  289,  292,  293 
detection,  249 

hydroxid,  139 

ions,  138 
Starch,    reaction   with   iodin,   210,    296, 

297 

Stibin,  166 
Stibnite,  167 
Strontium,  97 

detection,  270,  271 

flame,  97,  270 

ion,  97 

properties  of  metal,  97 

salts,   98 

spectrum,  /,  97 
Sublimation,  definition,  54 

reactions,  244 
Substances,  dissociating,  13 

inorganic,  13 

insoluble,  34 

insoluble,  how  indicated,  3 

in  solution,  condition  of,  9 

in  solution,  how  indicated,  3 

not  dissociating,  13 

organic,   13,   120-132 

pure,  4 

soluble,  34 
Sulfanilic  acid  mixture,  149,  302 

preparation  of  reagent,  320 
Sulfates,  188 

acid,  188 

action  of  heat,  187,  188 

detection,  278,  279,  290 
Sulfids,  180-184 

action  of  heat,  183 

detection,  183,  290,  294,  295 

oxidation,  180,  183,  184,  255 

poly-,  184 

precipitation,  252,  253,  268,  269,  290 

properties,   252 

solubility,  180-183,  252 
Sulfites,  185,  186 


action  of  heat,  185,  186 
detection,  290,  291 
oxidation,  185,  283 
reduction,  186 
separation,  291 
solubility,   186 
Sulfo-.     See  Thio- 
Sulfur,  179 

acids,  180,  184,  186,  188,  189 

as    oxidizing    agent,    125,    126,    139, 

162,  167,  188 

detection,  179,  180,  308,  309 
dioxid,  179 
flame,  179 
flowers,  179 
group,  172,  173 
hydroxids,  45 
in  alloys,  240 
ions,  180 
plastic,  179 
precipitation,  179,  184,  189,  225,  252, 

253,  255,  257,  259,  276,  277,  294, 

295 

properties  of  element,  179 
removal,  241 
trioxid,  179 
water,  180 
Sulfuric  acid,  186 
anhydrid,  179 
as  oxidizing  agent,  79,  84,  106,  124, 

183,  187,  245 
detection,  278,  279,  290 
fuming,  186 
heating  with   concentrated,  55,   123, 

124,  126,  130,  131,  183,  196,  199, 

203,  206,  209,  212,  245,  260,  278, 

279,  280,  298,  299,  300,  302,  303, 

304,   305 
ions,  187 

preparation  of  dilute  reagent,  320 
pyro,  186 

specific  gravity  of  solutions,  322 
Sulfurous  acid,  184,  185 
anhydrid,  179,  185 
detection,  290,  291 
as  oxidizing  agent,  185 
as  reducing  agent,  48,  185,  283,  291, 

293,  304,  305,  306,  307 
removal,  283 


350  INDEX. 


salts,  185,  186 
Symbols,  for  elements,  324 

meaning  of,  5 
Synthesis,  3 


Table  of  barium  salts,  101 

elements,  324,  326 

potassium  salts,  74 

sodium  salts,  71 

solubilities,  330,  331 
use  of,  243 

silver  salts,  88 

specific  gravities,  322,  323,  326 

spectra,  / 
Tantalum,  172 
Tartar  emetic,  167 
Tartaric  acid,  124 

as  reducing  agent,  47,  48,  124 

detection,  284,  285 

salts,  124,  125 
Tartrates,  124 

action  of  heat,  124 

detection,  284,  285 

solubility,   124 
Tellurium,  64,  172,  173 
Temperature  affects  reactions,  7 

affects  solubility,  32 
Tetrathionic  acid,  189 
Terbium,  65 
Tests,  borax-bead,  247 

drop,  246 

film,  246 

flame,  247 

match,  246 

on  charcoal,  245,  246 

special,  249 
Thallium,  118 

spectrum,  / 
Thenard's  blue,  115 
Thioantimonates,  169,  253 
Thioantimonites,  168,  253 
Thioarsenates,  164,  253,  259 
Thioarsenites,  162,  253 
Thiocyanates,  129 

as  reagent  for  cobalt,  231 

as  reagent  for  ferric  iron,  129,  227, 
231 

detection,  292,  293 


oxidation,  129,  291 

removal,  291 
Thiocyanic  acid,  129 

as  reducing  agent,  48,  291 
detection,  292,  293 
Thiostannates,  139,  141,  142,  253 
Thiostannites,  139,  253 
Thiosulfates,  189,  190 

action  of  heat,  189 

detection,  290,  291 

oxidation,  189,  290,  291 

separation,  291 

solubility,  189 
Thiosulfuric  acid,  188 

as  reducing  agent,  48,  290,  291 
detection,  290,  291 
ions,  190 
salts,  189,  190 
Thorium,  119 
Thulium,  65 
Tin,  137 

blowpipe  reaction,  138 

detection,  249,  256,  257 

group,  137 

ions,  138 

properties  of  metal,  137,  257 

salt,  140 

salts,  138-142 

Titanium,   119,   178,  306,  307 
Toning  solutions,  90 
Topaz,  133 
Transposition  with  sodium  carbonate,  68, 

69,  242,  260,  261 
Tungsten,  64,  190 
Turmeric  paper,  114,  282,  283 
Turnbull's  blue,  131,  223 


U 


Uranium,  64,  190 


Valence,  12 

high,   produces   strong   acids,  46 
how  indicated,  12 
increased  by  oxidation,  45 
increase  of  —  causes  increase  of  acid 

properties,  45 

low,  produces  strong  bases,  46 
of  elements,  324 


INDEX. 


351 


of  ions,  11 
Vanadium,  172 
Verdigris,  79 
Viscosity,  definition,  37 
Vitriol,  blue,  82 

green,  223 

white,  104 
Vitriols,  188 
Volatility,  as  determining  factor,  26 

W 
Water,  176 

ammonia,  147 

as  acid,  29,  176,  177 

as  base,  29,  177 

as  solvent,  13,  34,  177,  239 

baryta,  99 

bromin,  208 

chlorin,  200 

dissociation,  29,  176 

glass,  135 

hard,  93,  95,  96 

hydrogen   sulfid,   180 

iodin,  210 

preparation,  319 

lime,  95 

of  crystallization,  33 

sulfur,  180 

"Weathering"  of  rocks,  27 
Weight,  atomic,  57 

Apothecaries',  328 


English,  328 

equivalent,  58,  59 

Metric,  328 

molecular,   57 

Weights  and  measures,  57-60,  328 
"Wet  way"   reactions,  5 
Wood  alcohol,  114,  120,  282 


Xenon,  65 


Yellow  ammonium  sulfid  (see  ammonium 

poly  sulfid) 
Ytterbium,  64 
Yttrium,  111 

Z 

Zinc,  101 

action  with  alkalies,  102,  301 

blowpipe  reaction,  102 

detection,  268,  269 

dust,  101 

flame,  101 

group,  100,  101 

halids,  dissociation  of,  21 

hydroxid,  103 

properties  of  metal,  101,  102 

salts,  102-104 

Zincates,  102,  268,  269,  301 
Zirconium,  119 


y 


rBP 


14  DAY  USE 

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APR  08 1988 


021988 


DEC  05 1990 


AUTO  DISC  DEC  u  5 '9  ) 


LD  21A-50m-12,'60 
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University  of  California 

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